Classification of Elements and Periodicity in Properties

Classification of Elements and Periodicity in Properties (C) DOBEREINER TRIAD RULE : INTRODUCTION : The arrangement of all the known elements according (i) He made groups of three elements having similar chemical properties called TRIAD. to their properties in such a way that the elements of similar properties are grouped together in a tabular (ii) In Dobereiner triad, atomic weight of middle element is nearly equal to the average atomic weight of first form is called periodic table. and third element.Ex. (A) LAVOISIER CLASSIFICATION : Cl Br I (i) Lavoisier classified the elements simply in metals and 35.5 80.0 127 non metals. Metals are the one which have the tendency of losing the electrons. x= Na Na+ + e– Ca 40 and K K+ + e– Non-metals are the one which have the tendency of gaining the electrons. – F + e F – and – Cl + e Cl   Sr 87.5  x= Ba 137  – (ii) Drawback or Limitation : (a) As the no. of elements increases this classification become insufficient for the study of elements. Li Na K 7 23 39   x (b) There are few elements which have the properties of both metals as well as non-metals and they are Where x=average atomic weight called metalloids. Lavoisier could not decide where (iii) Other examples – (K, Rb, Cs), (P, As, Sb), to p lace the metalloids. (S, Se, Te), (H, F, Cl), (Sc, Y, La) 3 Classification of Elements and Periodicity in Properties (c) Metalloids (B, Si, As, Te, At etc.) and transition Key Points metals occupy bottom part of the curve. (a) For a dobereiner's triad, all the three elements should (d) Most electronegative i.e. halogens (F, Cl, Br, I) belong to the same group. occupy the ascending position on the curve. (b) For a dobereiner's triad, Difference in the atomic number should be 8 or 18. Drawback or Limitation : All the known elements Elements having similar properties occupy similar could not be arranged as triads. It is not applicable position on the curve. for d and f-block elements. Conclusion : On the basis of this curve Lother (D) NEWLAND OCTAVE RULE [1865] Meyer proposed that the physical properties of the (i) He arranged the elements in the increasing order of elements are periodic function of their atomic wt. their atomic mass and observe that properties of and this become the base of Mendeleef's periodic every 8th element was similar to the 1stone. like in table. the case of musical vowels notation (ii) At that time inert gases were not known. H Li Be B C N O F Na Mg Al Si P S Cl K Ca (iii) The properties of Li are similar to 8th element i.e. Na, Be are similar to Mg and so on. Drawback or Limitation : (F) TELLURIC HELIX : A three dimensional periodic (a) This rule is valid only upto Ca. because after Ca due table given by De-chan-chortois. to presence of d-block element there is difference (G) MENDELEEF'S PERIODIC TABLE of 18 elements instead of 8 element. (i) Mendeleef's periodic law : The physical and (b) After the discovery of Inert gas and included in the chemical properties of elements are the periodic period table it becomes the 8th element from Alkali function of their atomic weight metal so this law had to be dropped out. (ii) Characteristic of Mendeleef's periodic table : (E) LOTHER MEYER'S CURVE [1869] : (a) It is based on atomic weight (i) He plotted a curve between atomic weight and (b) 63 elements were known, noble gases were not atomic volume of different elements. discovered. (ii) The following observation can be made from the (c) He was the first scientist to classify the elements in a curve – systamatic manner i.e. in horizontal rows and in (a) Most electropositive elements i.e. alkali metals (Li, vertical columns. Na, K, Rb, Cs etc.) occupy the peak positions on (d) Horizontal rows are called periods and there were the curve. 7 periods in Mendeleev's Periodic table. (b) Less electropositive i.e. alkali earth metal (Be, Mg, Ca, Sr, Ba) occupy the descending position on the curve. (e) Vertical columns are called groups and there were 8 groups in Mendeleev's Periodic table. 4 Classification of Elements and Periodicity in Properties (f) Each group upto VIIth is divided into A & B – Corrections were done in atomic weight of elements subgroups.'A' sub groups element are called normal are – U, Be, In, Au, Pt. elements and 'B' sub groups elements are called (iv) Demerits of Mendeleef's periodic table : transition elements. (a) Position of hydrogen : Hydrogen resembles both, (g) The VIIIth group was consists of 9 elements in three the alkali metals (IA) and the halogens (VIIA) in rows (Transitional metals group). properties so Mendeleev could not decide where to place it. (h) The elements belonging to same group exhibit similar (b) Position of isotopes : As atomic wt. of isotopes properties. differs, they should have placed in different position (iii) Merits or advantages of Mendeleef's periodic in Mendeleef's periodic table. But there were no such table : places for isotopes in Mendeleev's table. (a) Study of elements : First time all known elements were classified in groups according to their similar (c) Anomalous pairs of elements : There were some pair of elements which didnot follow the increasing properties. So study of the properties of elements order of atomic weights. become easier . Ex. Ar and Co were placed before K and Ni respectively (b) Prediction of new elements : It gave in the periodic table, but having higher atomic encouragement to the discovery of new elements as weights. some gaps were left in it. Sc (Scandium) Ga (Gallium) Ge (Germanium) Tc (Technetium) These were the elements for whom position and (d) Like elements were placed in different groups: properties were well defined by Mendeleev even There were some elements like Platinum (Pt) and before their discoveries and he left the blank spaces Gold (Au) which have similar properties but were for them in his table. placed in different groups in Mendeleev's table. Ex. Blank space at atomic weight 72 in silicon group Pt Au was called Eka silicon (means properties like silicon) VIII IB and element discovered later was named Germanium. (e) Unlike elements were placed in same group: Similarly other elements discovered after mendeleef periodic table were. Eka aluminium - Galium (Ga) Eka Boron - Scandium (Sc) Li Eka Silicon - Germanium (Ge) Na Eka Mangense - Technetium (Tc) More reactive K Cu Less reactive (c) Correction of doubtful atomic weights : Rb Ag Coin metal Correction were done in atomic weight of some Alkali metal Normal elements Cs Au Transition elements. element Fr Initially, it was found that equivalent weight of Be is Cu, Ag and Au placed in Ist group along with Na, 4.5 and it is trivalent (V = 3), so the weight of Be K etc. While they differ in their properties (Only was 13.5 and there is no space in Mendeleev's table similar in having ns1 electronic configuration) for this element. So, after correction, it was found that Be is actually divalent (V = 2). So, the weight of (f) It was not clear that 'lanthanides and Actinides' are related with IIIA group or IIIB group. Be became 2 × 4.5 = 9 and there was a space between Li and B for this element in Mendeleev's (g) Cause of periodicity : Why physical & chemical table. properties repeated in a group. 5 Classification of Elements and Periodicity in Properties 6 Classification of Elements and Periodicity in Properties (H) MODERN PERIODIC TABLE (MODIFIED (v) Characteristics of modern periodic table : MENDELEEF PERIODIC TABLE) : (a) 9 vertical columns called groups. (i) It was proposed by Moseley. (ii) Modern periodic table is based on atomic number. (b) Ist to VIII group + 0 group of inert gases. (c) Inert gases were introduced in periodic table by (iii) Moseley did an experiment in which he bombarded Ramsay. high speed electron on different metal surfaces and obtained X-rays. (d) 7 horizontal rows called periods. He found out that   where = frequency of X-rays, Z = atomic number (I) LONG FORM / PRESENT FORM OF MODERN PERIODIC TABLE : from this experiment, Moseley concluded that the (It is also called as 'Bohr, Bury & Rang, Werner physical and chemical properties of the elements are Periodic Table) periodic function of their atomic number. It means that when the elements are arranged in the increasing (i) It is based on the Bohr-Bury electronic configuration order of their atomic number elements having similar concept and atomic number. properties gets repeated after a regular interval. This (ii) This model is proposed by Rang & Werner is also known as 'Modern Periodic Law'. (iii) 7 periods and 18 vertical columns (groups) (iv) Modern periodic law : The physical & chemical (iv) According to I. U. P. A. C. 18 vertical columns are properties of elements are the periodic function of named as Ist to 18th group. their atomic number. (v) The co-relation between the groups in long form of periodic table and in modern form of periodic table are given below – IA, IIA, IIIB, IVB, VB, VIB, VIIB, 1 2 3 4 5 6 7 8 IB, 9 10 11 IIB, IIIA, IVA, VA, VIA, VIIA, 0 12 13 14 15 16 17 18 7 Classification of Elements and Periodicity in Properties (vi) Elements belonging to same group having same number of electrons in the outermost shell so their properties are similar. 8 Classification of Elements and Periodicity in Properties (i) The group containing most electro negative elements (vii) Important Points : – GROUP VIIA (a) 2nd period elements (Li, Be, B) Shows diagonal relationship with 3rd period elements (Mg, Al, Si) (j) The group containing maximum number of gaseous elements–GROUP ZERO(18th) so (Li, Be, B) are called Bridge elements. Because of same ionic potentialvalue they shows similarity in (k) The group in which elements have generally ZERO valency – GROUP ZERO(18th) properties. (Ionic potential = Charge/Radius) (l) In the periodic table Number of Gaseous elements – 11(H, N, O, F, Cl + Noble gases) (b) 3rd period elements (Na, Mg, Al, Si, P, S, Cl) are Number of Liquid elements – 6 (Cs, Fr, Ga, Hg, called typical elements because they represent the Br, Uub) properties of other element of their respective group. Bromine is the only non-metal which exists in liquid (c) In 6th period all types of elements are included (s, p, form. d & f) Number of Solid elements – 95 (if discovered (d) No inert gas in 7th period. elements are 112) (e) Normal elements present in all periods. (m) No p-block elements in 1st and 7th periods. (f) Atomic number of last inert gas element is 86. (n) Ist period has all the elements in gaseous form (H, He) (g) Long form of modern periodic table can be divided (o) 0/18 group have all the elements in gaseous form. into four portions. (p) 2nd period contains maximum number of gaseous 1. Left portion – IA & IIA – s-block. elements. They are 4 (N, O, F, Ne) 2. Right portion – IIIA to VII A + 0 group – p(q) IIIB/3rd group is called longest group having 32 block. elements including 14 Lanthanides and 14 Actinides 3. Middle portion – IIIB to VIIB + VIII + IB IIB – Sc d-block. Y 4. Bottom portion – IIIB – f-block elements La................Lanthanides (14) (h) The group containing most electro positive elements Ac................Actinides (14) – GROUP IA. (viii) Nomenclature of elements : (a) IUPAC gave names to elements above atomic number 100 as follows – 0 1 2 3 4 5 6 7 8 nil un bi tri quad pent hex sept oct 9 enn (b) In all the elements suffix is – ium. Ex. Atomic No. IUPAC Name 101 Un nil unium 102 Un nil bium 103 Un nil trium 104 Un nil quadium 105 Un nil pentium 106 Un nil hexium 107 Un nil septium 108 Un nil octium 109 Un nil ennium 110 Un un nilium Symbol Md No Lr Rf Db Sg Bh Hs Mt Ds Symbol Unu Unb Unt Unq Unp Unh Uns Uno Une Uun Elemental Name Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassnium Meitnerium Darmstadtium 9 Classification of Elements and Periodicity in Properties (ix) Merits of long form of periodic table : (b) Their general electronic configuration is : (a) Position of isotopes – Atomic number of isotopes IA IIA are similar, so different isotopes can be placed at ns1 ns2 same place in periodic table. (b) (Ar – K) (Co – Ni) (Te – I) are now in increasing order of atomic number. (c) Lanthanides and actinides are in IIIB group. IIIA IVA VA VIA VIIA (d) In modern periodic table diagonal line separates out ns2 np1 ns2 np2 ns2 np3 ns2 np4 ns2 np5 metals, metalloids and non metals. (e) Elements of same group have same general formula of electronic configuration of outer most shell. (c) These elements lies in group IA to VIIA and st th (x) Demerits of long form of periodic table : period I to 7 (a) Position of hydrogen is still controversial. (d) Elements of 2nd period known as bridge elements. (b) 'He' is a inert gas but it has different electronic (e) Elements of 3rd period (Na, Mg, Al, Si, P, S and Cl) configuration than other inert gas elements. are called typical elements These are 7 in numbers. (c) Lanthanides and actinides are still not placed in main (iii) Transition elements : frame. Old concept (Given by Mendeleev's) : They are (d) Isotopes have different physical properties but have called transition elements because their properties same place in periodic table. are intermediate of s and p- block elements or they CLASSIFICATION OF ELEMENTS transit the properties of metals to non metals. (A) Bohr classification : Elements can be divided in four Note : According to this concept, Zn, Cd, Hg and Uub parts on the basis of electronic configuration. are transition elements. (i) Inert gas elements : New concept (Given by Bohr) : The elements in (a) The elements in which Ultimate orbit is completely which both ultimate (n) as well as penultimate shells filled are called inert gas elements. (n – 1) are incomplete either in atomic state or in (b) General electronic configuration is ns2p6 (Except He some oxidation state are called transition elements. = 1s2) Note : According to this concept, Zn, Cd, Hg and Uub (c) Because of most stable configuration, they are very are not transition elements because they do not have less reactive. Hence known as noble gas or inert incomplete penultimate shell either in atomic state or gas element. in some oxidation state. (d) These elements are present in '0' group or 18th group Zn = [Ar] 3d10 4s2 and Zn+2 = [Ar] 3d10 and 1 to 6th period of periodic table. (a) group – IIIB to VIIB + VIII + IB and (e) Number of inert gas elements are 6(One in each Periods – 4th to 7th period upto 6th) He, Ne, Ar, Kr, Xe, Rn (ii) Normal or Representative elements : (b)Electronic configuration (n–1)d1–10 ns1 or 2 (a) The elements in which ultimate orbit is incomplete, while penultimate orbits are complete are called normal elements. (c) Total number of d-block elements = 40 Total number of transition elements = 36 (Except Zn, Cd, Hg and Uub) Note : All transition elements are d-block elements but all d-block elements are not transition elements. 10 Classification of Elements and Periodicity in Properties (d) Total s - block elements are (14) (including H and (iv) Inner transition elements : He) (a) The Elements in which all the three shells that is ultimate (n) penultimate (n–1) and pre or (e) Total s - block metals are (12) (excluding H), H is non-metal antipenultimate (n–2) shell are incomplete are called General characteristics of s-block elements inner transition elements. Ex. 58Ce = [Xe] 6s2, 5d1, 4f1  s-block elements are soft metal, that is why these elements have low melting and boiling point.  Their oxides and hydroxides are basic in nature.  They act as reducing agent.  They have very low ionisation energy and highly electropositive so, they form ionic compound.  Flame of these elements have the property of showing different colour. (b) Electronic configuration (n – 2)f1 – 14 (n – 1) d0 or 1 ns2 (c) These are 28 in number. (d) Group – IIIB and Period – 6th & 7th (e) Inner transition elements are divided into two series. (I) Lanthanide series/Rare earth elements / Lanthenones : (Ce58 – Lu71,14 elements) The first element of this series is Cerium & not Lanthanum. In these elements, last electron enters into 4f subshell They are present in IIIB group and 6th period of the periodic table. Promethium (61Pm) is the only lanthanide which is synthetic and radioactive in nature. (II) Actinide series /Man made elements /Actinones (Th90 – Lw103 14 elements) (ii) p-block elements : (a) Last e– enters in p - sub shell (b) Group – IIIA (13) to VIIA (17) + 0 group (18) (except He) and Period – 2nd to 6th (c) Electronic configuration – ns2 np1 – 6 and - block elements – (30) Total p Characteristic of p-block elements  p-block possesses all three kind of elements i.e. metals, non-metals and metalloids. The first element of this series is Thorium & not Actinium because last electron enters into 5f subshell.  Oxides of non-metals are acidic in nature. They are present in IIIB group and 7th period of the periodic table.  They form covalent compounds mostly.  They are oxidising in nature. All the actinides are radioactive in nature. First three (iii) d-block elements : – elements (Th, Pa, U) are found in nature while others (a) Last e enters in (n – 1)d subshell are synthetic in nature.After U92 i.e. from 93Np (b) Group – IIIB – VIIB, VIII, IB, IIB or group 3 to 12 (IUPAC) and Period – 4th to 7th onwards elements are called transuranic elements They are heavier than uranium and they are derived (c) Electronic configuration – (n –1)d1 – 10 ns1 or 2 from uranium by nuclear reactions.   (B) Classification on the Basis of Sub shell in which (d)  last e– enters : (i) s-block elements : (a) In these elements last electron enters in s - subshell. (e) IIB elements (Zn, Cd, Hg, Uub) are d-block elements but not transition elements. (b) Group – IA, IIA + 0 group (He only) and Period th – 1 to 7 (c) Electronic configuration – ns1 – 2 IA group n = 1 to 7 and IIA group n = 2 to 7 11 Classification of Elements and Periodicity in Properties (iv) f-block elements : General characteristics of d-block elements (a) Differentiating e– enters in (n – 2) f subshell.  They all are metals, which are very hard having (b) Group – IIIB and Period – 6th and 7th high melting and boiling point. (c) From atomic number 58 - 71, 6 th period ;  Elements of this block situated in between 's' and Lanthanide series 4f1 – 14 5d0 or 1 6s2 p-block elements. So the character of this block 90 – 103, l 7th period ; Actinide series 5f1 – 14 elements lie between s and p-block elements 6d0 or 1 7s2  They show variable oxidation - state. (d) Total number of f block elements – (28) and +2 +3 +4 +5 +6 +7 All the actinides are radioactive elements. Ex. Mn = Mn , Mn , Mn , Mn , Mn , Mn (e) Transuranic actinides are man made elements.  They form lonic and covalent bond both. (Np93 – Lw103)  They are good conductor of heat and electricity (f) Lanthanides are found rarely on earth so these are and form complex compounds. called rare earth metals.  Metals, which have unpaired electrons show General characteristics of f-block elements – paramagnetism.  All the f-block elements are heavy metals.  They form alloys and most of the elements act as  It shows high melting and boiling point. catalyst.  The most common oxidation state of these elements is +3. [RPMT 2004]  The elements of 5f series are radioactive in nature. Predicting atomic number of the succesive member in a group of family (i) Magic Numbers : Knowing the atomic number of the first member of a group, we can write the atomic number of the subsequent elements by adding given magic number Ex. Group 1 2 3 4, 5, 6, 7, 8, 9 10, 11, 12, IA IIA IIIB 13, 14, IIIA IV 15, 16, 17 18 V 'O'gp. VI VII I+ 2 – 8 II+ 8 8 8 8 III+ 8 8 18 18 IV+ 18 18 18 18 18 18 V+ 18 18 18 32 32 32 VI+ 32 32 32 32 – – VII (ii) In group IA – Atomic number of H is 1 and atomic number of other element will be as follows – Magic number                  12 Classification of Elements and Periodicity in Properties Determination of period, block and group of an element (i) Period number: The period no. of the element can be predicted from the principal quantum no. (n) of the valence shell. (ii) Block number: Last electron enter in which orbital is knows as block no. (iii) Group number: It is predicted from the number of electrons in the valence shell and penultimate shell . Ex. S.No. Electronic Configuration 2 10 6 1 Period Block Prediction Group number number of Group number 5 s No. of ns electrons 1 1. [Ar]4s 3d 4p , 5s 2. [Kr] 5s2, 4d10 5p2 5 p ns ele. +10 +np ele. 2+10+2 =14 3. [Rn] 7s2, 6d4 5f14 7 d ns ele. + (n–1)d ele. 2+4 6 f — 4. 2 1 12 [Xe] 6s ,5d , 4f =6 3/III B PERIODICITY Old concept : Given by : Frankland Valency with respect to Hydrogen and Chlorine : Valency of H = 1 and Cl = 1. It is defined as the number of hydrogen or Chlorine atoms attached with a particular element. (a) In a period, the ultimate orbit remain same, but the IA IIA IIIA IVA VA VIA VIIA number of e– gradually increases. NaH MgH2 AlH3 SiH4 PH3 H2S H–Cl (b) In a group, the number of e– in the ultimate orbit NaCl MgCl2 AlCl3 SiCl4 PCl3 SCl2 C–Cl remains same, but the values of n increases. Valency 1 2 3 4 3 2 1 (ii) Causes of periodicity : Note : (i) The regular gradation in properties from top to bottom in a group and from left to right in a period is called periodicity in properties. (a) The cause of periodicity in properties is due to the same outermost shell electronic configuration coming at regular intervals. Valency w.r.t. H across the period increases upto 4 and then again decreases to 1. Valency with respect to oxygen : Valency of 'O' = 2 (b) In the periodic table, elements with similar properties It is defined as twice the number of oxygen atoms attached with a particular atom. occur at intervals of 2, 8, 8, 18, 18 and 32. These IA IIA IIIA IVA VA VIA VIIA numbers are called as magic numbers. Na 2O MgO Al2O3 SiO2 P2O 5 SO 3 Cl2O7 Periodic Properties Valency 1 2 3 4 5 6 7 Valency : It is defined as the combining capacity of the elements. The word valency is derived from an Note : Valency with respect to oxygen increases from 1 to Italian word "Valentia" which means combining 7 across the period. Valency w.r.t. 'O' is equal to the capacity. group number. 13 Classification of Elements and Periodicity in Properties New concept : This concept is based on the electronic configuration. According to this concept valency for IA to IVA group elements is equal to number of valence shell e– and from VA to zero group, it is – [8– (number of valence e–)]. Valency = No. of valence e– IA ns1 1 1 Valence shell e– Valency IIA ns2 2 2 IIIA ns2np1 3 3 IVA ns2np2 4 4 Note : All the elements of a group have same valencies because they have same number of valence shell electrons. Valency=(8–no.of valence e–) VA ns2np3 5 3 (8 – 5) = 3 VIA ns2np4 6 2 VII ns2np5 7 1 0 ns2np2 8 0 (8 – 8) = 0 (viii) Highest density of solid metal Ir (22.63 gm/cc) and Os (22.6 gm/cc) and liquid metal Hg(13.6 gm/cc) (ix) Order of density of elements : Li < K < Na < Rb < Cs Sc < Y < La Density : Ca < Mg < Be < Sr < Ba Ti < Zr << Hf (i) Density increases down the goup because increase He < Ar < Kr < Rn Zn < Cd < Hg in atomic mass is much greater than volume. (ii) Factors affecting density are Hg < Au < Os < Ir Sc < Zn < Fe < Hg (a) Inner shell configuration (b) Atomic weight (c) Packing capacity (d) Chemical bond (iii) In period : Density first increases till maximum and then decreases. (s-block to d-block increases, dblock to p-block decreases) (i) (ii) (iv) In group : Ex. In VIIA group – F and Cl are gases (Low density) Br is liquid (density 3.19 g/cm–3) (a) I is solid (density 4.94 g/cm–3) (vi) Exceptions – The density of sodium and magnesium is greater than K and Ca respectively. This can be explained by inner shell configuration. Density Na > K (b) Cr < Fe < Cu Fe < Cu < Pb Fe < Hg < Au Boiling points and melting points : It is a property of aggregate of atoms and not of a single atom. i.e. why it is a molecular propertry. Melting point of a solid depend on – (a) Structure of solid (b) Chemical bond (c) Bond energy Structure of solid made up of (i) atomic solid (ii) molecular solid (iii) metallic solid. Order of decreasing melting point is : Atomic solid > Metallic solid > Molecular solid Chemical bond among particles in solid. Order of bond strength is – Covalent bond > Metallic bond > Vander waal bond In K, 3d0 orbital is vacant, 3rd orbit has the capacity (c) Bond energy – Covalent solids like diamond, SiO 2 to hold 18e– but it has only 8e–. Due to empty orbitals etc have only covalent bonds between atoms so their atomic volumes increases and so density decreases bond energy is higher than molecular or metallic solids. (vii) Li is the lightest metal known. 14 Classification of Elements and Periodicity in Properties (iii) In period : Along the period from left to right boiling Ex. I2. point and melting point first increases then decreases. (v) Order of melting point : Li > Na > K > Cs C > Si > Ge Alkali metals – Crystal structure BCC (low boiling Be > Ca > Sr > Ba > Mg C > B > Si point & melting point) F2 < Cl2 < Br2 < I2 B > Al > Ga Transition metals – ............"..... FCC (High boiling point & melting point) He < Ne < Ar < Rn Inert gases ............ Lowest boiling point & Zn > Cd > Hg melting point (Vander waal force) Cu > Au > Ag Transition elements ........... Highest melting point Screening effect (  ) and effective nuclear Metals W(Tungston) maximum melting point (3410°C) charge (Zeff) : Hg(Mercury) lowest melting point (-38°C) (i) Valence shell e– suffer force of attraction due to nucleus. and force of repulsion due to inner shell Non metals electrons. Carbon (In the form of diamond) Highest melting (ii) The decrease in force of attraction on valence e– point (3727°C) due to inner shell e– is called screening effect or Helium Lowest melting point (-270°C) shielding effect.(i.e. total repulsive force is called shielding effect.) (iv) In Group : – (a) In s-block elements boiling point & melting point (iii) Due to screening effect valence shell e experiences less force of attraction exerted by nucleus.(i.e. total decreases down the group. attraction force experienced by valence e– is called Li, Na (solid) Cs, Fr (liquid) Zeff.) (b) In d-block elements boiling point & melting point (iv) There is a reduction in nuclear charge due to screening effect. Reduced nuclear charge is called effective increases down the group (due to lanthanide nuclear charge. contraction, zeff increases and hence bond energy (v) If nuclear charge = Z, Then effective nuclear charge increases) = Z – [Where (Sigma)= Screening constant] (c) In p-block elements So, Zeff = (Z – ) (i) From IIIA – IVA group boiling point & melting point Slater's rule to know screening constant () : decreases down the group and from VA to '0' group, – boiling point & melting point increases down the group. (a) Screening effect (S.E.) of one e of the 1s is 0.30 (Atomic or molecular weight vander waal force) (ii) boiling point and melting point of monoatomic molecules are lesser than diatomic molecules. '0' group < Halogens (iii) Atomic solid non metals like B, C and Si has higher boiling point and melting point due to strong covalent bond. (iv) Boiling point & melting point of molecular solids are less because of weaker vander waal force among molecules Ex. In He (1s2) Screening effect of one 1s e–. where = 0.30  Zeff = Z – = 2 – 0.30 = 1.7 (b) Screening effect of ns and np (Outermost orbit) electron is 0.35 (c) Screening effect of (n – 1) penultimate orbit s, p, d electrons is 0.85 (d) Screening effect of (n – 2) and below all the e– present in s, p, d, f is 1.0 15 Classification of Elements and Periodicity in Properties Ex. 1s22s1 (Z = 3) Lithium Potassium (Z = 19) Zeff Zeff on 2s electron = Z– = 3 – (2 × 0.85) = 3 – 1.70 = 1.30 Note: Only two electron present in (n–1) shell affects Z. Zeff Zeff 1 =Z– = 19 – ( 10 × 1 + 8 × .85) = 19 – 16.8 = 2.2 Sodium (Z = 11) Zeff = 11 – (2 × 1 + 8 × .85) = 11 – 8.8 = 2.2 (Effective Nuclear charge of elements of second period) Element Electronic Z Configaration 3 2 1 Li 1s 2s Be4 1s2, 2s2 of ns & np electron (n–1) Total Effective orbital Screeing nuclear Constant charge (a) (b) (a + b) Z*=Z –  3 – 0.85× 2=1.70 1.70 1.30 4 1× 0.35=0.35 0.85× 2=1.70 2.05 1.95 B 1s ,2s ,2p 1 5 2× 0.35=0.70 0.85× 2=1.70 2.40 2.60 C6 1s2,2s2,2p1 6 3× 0.35=1.05 0.85× 2=1.70 2.75 3.25 N7 1s2,2s2,2p3 5 O 8 F9 2 2 7 4× 0.35=1.40 0.85× 2=1.70 3.10 3.90 4 8 5× 0.35=1.75 0.85× 2=1.70 3.45 4.55 1s2,2s2,2p5 9 6× 0.35=2.10 0.85× 2=1.70 3.80 5.20 2 2 1s ,2s ,2p Periodicity : (c) That is why in a period Zeff increases by 0.65 and hence atomic size decreases considerably. (i) From left to right in a period Zeff increases (a) For normal / Representative elements, In a period (d) In transition series Z increase by + 1 but screening effect increases by 0.85 So Zeff is 0.15 nuclear charge increases by +1, while orbit remains same. (1– 0.85 = 0.15) [Because e– enters in (n – 1) orbit which has the value of = 0.85] (b) In a period e– enters in outermost shell and screening effect of outermost electron is 0.35. So by increasing In transition series Zeff increases by very less amount, 1 nuclear charge, Zeff increases by 0.65 (1 – 0.35 i.e. by 0.15 from left to right and hence atomic size = 0.65). remains almost constant. Element Zeff Sc 3.00 Ti 3.15 V 3.30 Cr 2.95 Mn 3.60 Fe 3.75 Co 3.90 Ni 4.05 Cu 3.70 Zn 4.35 (ii) From top to bottom in a group Zeff remain constant Element Li Na K Rb Cs Fr Zeff 1.30 2.20 2.20 2.20 2.20 2.20 16 Classification of Elements and Periodicity in Properties Case II ATOMIC RADIUS Heteroatomic molecule with more The average distance of valence shell e– from nucleus is called atomic radius. It is very difficult to measure electronegativity difference : the atomic radius because – Schoemaker and Stevenson law : (i) The isolation of single atom is very difficult. If in a diatomic molecule electronegativities of A – B (ii) There is no well defined boundary for the atom. have more difference. Then actual bond length will (The probability of finding the e– is 0 only at infinity). be reduced. So, the more accurate definition of atomic radius is– As per schoemaker & Stevenson– The reduction in  Half the inter-nuclear distance(d) between two bond length depends on the difference in atoms in a homoatomic molecule is known as electronegativities of atoms by following manner atomic radius. dA – B = rA + rB – 0.09 (XA – XB)  This inter-nuclear distance is also known as bond Here XA is E.N. of A & XB is E.N. of B length.Inter-nuclear distance depends upon the type of bond by which two atoms combine. Based on the chemical bonds, atomic radius is divided into four categories – Ex.1 If bond length of F2 = 1.44 Å, Bond length of H2 (A) Covalent radius (B) Ionic radius = 0.74Å. Find out the bond length of H – F ? (C) Metallic radius (D) Vander waal radius (EN of F is 4.0, EN of H is 2.1) (A) Covalent radius (SBCR –Single Bonded Covalent Sol. d = rF + rH – 0.09 (XF – XH) H –F Radius) dF–F = 1.44 / 2 = 0.72 Å, dH – H = 0.72/2 = 0.37 Å (a) Covalent bonds are formed by overlapping of atomic  dH–F = 0.72 + 0.37 – 0.09 ( 4.0 – 2.1) orbitals. = 1.09 – (0.09 × 1.9) = 1.09 – 0.171 = 0.919Å (b) Internuclear distance is minimum in this case. (c) Covalent radius is the half of the internuclear distance (B) Ionic Radius between two singly bonded homo atoms. Ex. If internuclear distance of A–A(A2) molecule is (i) Cationic radius : (a) When an neutral atom loses e– it converts into cation (dA – A) and covalent radius is rA then (+ve charged ion) dA–A = rA + rA or 2rA   (b) Cationic radius is always smaller than atomic radius Ex. In Cl2 molecule, internuclear distance is 1.98A0 because after loosing e– number of e– reduces, but number of protons remains same, due to its Zeff  so rcl = increases, hence electrons pulls towards nucleus and (d) SBCR of O, N and C etc. elements can be atomic radius decreases, moreover after loosing all determined by taking H 2 O 2 , N 2 H 4, C 2 H 6 the electrons from the outer most shell, penultimate respectively. shell becomes ultimate shell which is nearer to nucleus Case I so size decreases. For Heteroatomic molecule with less electronegativity difference. (c) Size of cation  For A – B molecule electronegativities of A & B are Ex. (i) Fe > Fe+2 > Fe+3 approximately equal Ex. C – I (ii) Pb+2 > Pb+4 Electronegativity of C & I are approx equal (2.5) (iii) Mn > Mn+2 > Mn+3 > Mn+4 > Mn+5 > Mn+6 > Mn+7 internuclear distance of C – I is 2.13Å (ii) Anionic radius dC–I = rC + rI, ( rC is 0.77 Å) (a) When neutral atom gains e– it converts into anion  rI = 2.13 – 0.77 = 1.36Å 17 Classification of Elements and Periodicity in Properties (b) Anionic radius is always greater than atomic radius (D) Vander Waal’s radius because In an anion e– are more than protons so (a) Those atoms (like noble gases) which are not bonded effective nuclear charge reduces, and inter electronic with each other, experiences a weak attractive force repulsion increases, which also increases screening to come nearer. effect. So distance between e– and nucleus increases (b) Half of the distance between the nuclei of adjacently and size of anion also increases. placed atoms in solid state of a noble gas is Vander Ex. Flourine (Z=9) Waal's radius. F F– (c) Vander Waal radius > Covalent radius. Proton 9 9 (d) Inert gas have only Vander Waal radius. e– 9 10 (e) In molecules of nonmetals solid both covalent and Vander Waal radius exists. so = =1 = 0.9 As Zeff of F– is less than F so size of F– > F (c) Size of iso electronic species :  Those species having same number of e– but different nuclear charge forms isoelectronic series.  For isoelectronic species the atomic radius increases with decrease in nuclear charge Species K+ Ca+2 Ar S–2 Cl– Z 19 20 18 16 17 e 18 18 18 18 18 Vander Waal radius = 2 × Covalent radius (f) Vander Waal radius > Metallic radius > Covalent radius (g) Vander Waal force of attraction  Molecular weight or atomic weight (in inert gases) (h) In a period from left to right Vander Waal radius decreases. (i) In a group of non metal from top to bottom its values increases. Factors affecting atomic size are : Order of radius : (S–2 > Cl– > Ar > K+ > Ca+2), (N3–> O2–> F–> Ne > Na+ > Mg+2 > Al+3) (C) Metallic radius (a) Half of the nuclear distance between two adjacent metallic atoms in crystalline lattice structure. (b) there is no overlapping of atomic orbitals so, Metallic (a) radius > Covalent radius < (c) Metallic radius  (d) More metallic radius  less bond strength Li > Be > B > C > N > O > F (b) Atomic radius number of shells loose crystal packing  (body centred packing) (e) Less metallic radius  Tight crystal packing FCC High bond strength (Hexagonal close packing)  Li < Na < K < Rb < Cs (c) Atomic radius  Screening effect (d) Atomic size Magnitude of –ve charge O < O– < O–2 (e) Atomic radius  Mn > Mn+2 > Mn+3 > Mn+4 18 (f) (i) Ex. (ii) Ex. Classification of Elements and Periodicity in Properties (a) At. size of Ga = At. size of Al, due to transition Atomic radius  contraction. (b) In transition elements nuclear charge increases by 1. — > —N N— >N N (c) but e– enters in (n –1)d orbital exerts screening effect. Periodic variation of atomic size : (d) Screening effect of (n –1)d e– balance the nuclear Across a period : It decreases from left to right in a charge by 85% period as nuclear charge increases (e) Zeff on increasing each electron = 1 – 0.85 = 0.15 Li > Be > B > C > N > O > F < Ne (f) Increase in nuclear charge is only 0.15 so atomic In a group : It increases from top to bottom in a size remains almost constant. group as number of shell increases Li < Na < K < Rb < Cs Different orders of atomic and ionic radii : Exceptions : Transition elements Na+ < Ne < F— Ni < Cu < Zn           Order of atomic radius – Sc > Ti > V >Cr Mn > Fe Co Ni < Cu < Zn (a) (b) (c) (d) (e) Be < Li < Na Cr < V < Ti < Sc Cl < Na < Ar Cu < Au < Ag H+ < H < H— Ti < Zr Hf + — I < I < I Sc < Y < La Lanthanoide Contraction : IONISATION POTENTIAL OR Outermost electronic configuration of inner transition IONIZATION ENERGY OR elements is IONIZATION ENTHALPY (n –2) f1 – 14, (n–1)s2p6d0–1, ns2 (n = 6 or 7) (i) Minimum energy required to remove most loosly e– enters in (n – 2) f orbitals bonded outer most shell e– in ground state from an Mutual screening effect of e– is very less, because of isolated gaseous atom is known as ionisation complicated structure of f-orbital potential. Nuclear charge increases by one (+1) in lanthanoides (Isolated  Without any bonding with other atom) and actinides so atomic size of these elements slightly decreases. It is known as lanthanoide contraction. (ii) Successive Ionisation Energy : Here Nuclear charge > Screening effect. (a) For an atom M, successive ionisation energies are st nd rd as follows In I , 2 and 3 transition series, Radii– 3d < rd 4d 5d (except III B) M + E1  M+ + e– IIIB IVB E1 = Ionisation Potential size Sc Ti  size increases M + + E2  M+2 + e– increases Y Zr Equal due to La Hf E2 = Ionisation Potential lanthanide contraction M+2 + E3  M+3 + e– Transition contraction : IIIA B Fe+2 > Fe  (iv) In a group : Size increase so ionisation potential (C) Screening effect : Higher is the screening effect on decrease. the outer most electrons causes less attraction from Li Na K Rb Cs the nucleus and can be easily removed, which is  leading to the lower value of Ionisation Potential   (D) Penetration power of sub shells : (a) Order of attraction of subshells towards nucleus (Penetration power) is s > p > d > f will be required to remove e– from s-subshell in comparison to p,d & f. Ex. Be B 1s2, 2s2 1s2, 2s2 2p1 Ionisation Potential Be > B – After loosing one e , B attains electronic configuration of Be, so Be. ionisation potential of B is more than Ionisation Potential of B > Be Size increases, Ionisation Potential decreases Exception : Ionisation Potential Ga > Al (While Ionisation Potential decreases down the group it is due to Transition contraction)  (v) In a period : In a period atomic size decreases and zeff increases so removal of electron becomes difficult and ionisation potential increases. Li Be B C N O F Ne  atomic size decreases, zeff increases, Ionisation Potential increases. Order of ionisation potential Li < B < Be < C < O < N < Fe < Ne 20 Classification of Elements and Periodicity in Properties (vi) Application of ionisation potential : Ex. Mg  Mg+ Ist ionisation potential 7.4 eV Mg+  Mg+2 IInd ionisation potential } (A) Metallic and non metallic character : Difference of ionisation potential < 11 eV So Mg+2 Metallic  Ionisation Potential Low is more stable. (Na, K, Rb etc.) Al  Al+ non metallic  Ionisation Potential High + } 12.8 eV So Al is more Al+  Al+2 (F, Cl, Br etc.) +3 } 6.0 eV So Al is more stable Al+2  Al+3  Al+ is stable only in gaseous state {Al+3 is stable in liquid and solid state. (B) Reactivity : (E) Basic nature : It is property of elements with loosely held electrones   Reactivity increases down the group as ionisation potential decreases. (C) Reducing character : M2O > MO , Na2O > MgO , NaOH > Mg(OH)2 Cs2O > Rb2O > K2O > Na2O > Li2O (F) To determine the number of valence electron :  Ex. Li Ist Ionisation IInd Ionisation (a) IA group has minimum ionisation potential so they Potential Potential are strong reducing agents in gaseous state 5.4 e V 75 eV ( Li < Na < K < Rb < Cs) High energy difference in two successive ionisation potential shows change in energy levels ie. valence (b) IA group – In Aqueous state shell has only one electron. reducing character Li > K ~ Rb > Cs > Na As the degree of hydration is more in Li due to high (G) Ionic nature : Ionic nature of elements decreases from left to right in a period. charge density. (D) Stability of oxidation states :  (a) If the difference between two successive ionisation potential > 16eV then lower oxidation state is stable. (H) Nature of hydroxide : In a hydroxide AOH – Nature of hydroxide depends Ex. Na  Na+ Ist ionisation potential 42.7 eV on ionisation potential of A. } Na+  Na+2 IInd ionisation potential A–O–H . If ionisation potential of A is lower than Difference of ionisation potential > 16 eV So Na+ oxygen, it will loose electron and will be basic is more stable. A–OH  A+ + OH– (b) If the difference between two successive ionisation If A has high ionisation potential than oxygen – potential < 11 higher oxidation state is stable. A – O – H  AO– + H+ then the hydroxide will be acidic in nature. NaOH Mg(OH)2 Al(OH)3 Si(OH)4 P(OH)3 SO2(OH)2 ClOH  Ionisation Potential increases Acidic nature increases Basic nature decreases Note : Ionisation potential decreases in a group so, basic nature increases. 21 Classification of Elements and Periodicity in Properties ELECTRON AFFINITY/ELECTRON GAIN ENTHALPY (i) The amount of energy released or absorbed when electron is added to the valence shell of an isolated gaseous atom known as Electron affinity. X + e– Electron Affinity  X– + Ex.       (v) On adding 2nd electron in anion X– X– + e–  X–2 – IInd Electron Affinity or X– + e– + IInd Electron Affinity  X–2 eg H (2nd electron gain enthalpy) = + ve (iii) If Ist Electron affinity is exothermic (vi) IInd Electron affinity is endothermic process, i.e. energy has to given to introduce 2nd electron in an anion, because of repulsion between negative charge on anion (X–) and 2nd electron.IInd Electron Affinity is always negative. eg H (electron gain enthalpy) Ex. O(gas) + e–O– (gas) + 168 K.cal (exothermic) (ii) Mostly energy is released in the process of first Electron affinity. X + e–  X– + Ist Electron Affinity = – ve st I Electron affinity is mostly positive (iv) Electron affinity of neutral atom is equal to ionisation potential of its anion. O– + e– O–2 – 912.7 K.cal Net reaction O(g) + 2e–   O–2(g) Electron Affinity = – 744.7 K.cal ( endothermic) {Energy evolved in Electron affinity to convert neutral atom to anion}={Energy absorbed to remove e– from anion} (vii) Factors affecting electron affinity : (A) Atomic size (B) Screening effect (C) Effective nuclear charge (Zeff)   Electron Affinity  Zeff (D) Stability of completely filled or half filled (viii) In period : orbitals : (a) Electron affinity increases along the period due to Electron affinity of filled or half filled orbital is very increase in zeff and decrease in atomic size. less or zero. (b) Halogens possess maximum electron affinity due to Ex. small size and maximum zeff and after gaining one electron. They attain a stable inert gas configuration. (i) Elements (He, Ne, Ar , ..), Electron Affinity = Zero (ii) Elements (Be, Mg, Ca, ....), Electron Affinity = Approximately Zero Li B C O F Na Al Si S Cl zero Less zero (iii) Elements (N, P, As, .... ), Electron Affinity = Very less In 2nd period – Ne < Be < N < B < C < O < F In 3rd period – Ar < Mg < P < Al < Si < S < Cl 22 Classification of Elements and Periodicity in Properties (ix) In Group : (a) Electron affinity decreases down the group. As the size increase and attraction on additional electron is Note : less. N & P have low electron affinity due to stable half (b) Electron affinity of 3rd period element is greater than filled configuration. electron affinity of 2nd period elements of the respective group. (x) Application of Electron affinity : F Cl 1s2, 2s22p5 [Ne] 3s23p5 Due to small size of fluorine, electron density around the nucleus increases. The incoming electron suffers more repulsion. In case of chlorine electron density decreases due to large size, decreasing order of electron affinity (a) Electron affinity  Oxidising nature But F has more oxidising power than Cl because F. (i) has high standard reduction potential. (ii) Has high hydration enthalpy. (iii)Low bond dissociation energy. (xi) Difference between electronegavity and Electron Affinity : Electronegativity Electron Affinity Tendency of an atom in a molecule Energy released when an electron is added to to attract the bonded electrons neutral isolated gaseous atom It regularly changed in a period It does not changes regularly It has no unit It is measured in eV/ atom or KJ mol–1 or K cal mole–1 ELECTRONEGATIVITY (EN) (i) The tendency of an atom to attract shared electrons towards itself is called electronegativity. or (b) A – B  A+ + :B– (Electronegativity A < Electronegativity B) depending on their tendency to attract bonded electron. (ii) Electronegativity and Electron affinity both have tendency to attract electrons but electron affinity is (iv) There is no unit of electronegativity as for isolated atoms. Where as electronegativity is for electronegativity is tendency of a bonded atom not bonded atoms. an energy (iii) A polar covalent or ionic bond of A – B may be broken as – + (a) A – B  A : + B (Electronegativity A > Electronegativity B) (v) Pauling explained it first time. (vi) Electronegativity of some other elements are as follows – 23 Classification of Elements and Periodicity in Properties Note : Small atoms are normally having more electronegativity than larger atoms. (vii) Factors Affecting electronegativity : (A) Atomic size (B) Effective nuclear charge (Zeff)  (C) Hybridisation state of an atom  (D) Oxidation state  2 3 sp > sp > sp s character 50% 33% 25% Electronegativity 3.25 2.75 2.5 Key Points :   s-orbital is nearer to nucleus so by increasing scharacter in hybridisation state, Electronegativity also increases. Electronegativity does not depends on filled or half filled orbitals, because it is a tendency to attract bonded electron, not to gain electron from out side. (viii) Periodic table & Electronegativity : (a) Electronegativity decreases down the group. (b) In period on moving from left to right electronegativity increases. Exceptions : In '0' group – Electronegativity is always zero, because inert gas do not form molecule.  +2 +4 Mn < Mn < Mn+7 O –2 < O –1 < O < O+1 < O+2 Fe < Fe+2 < Fe+3 (c) Electronegativity of Cs and Fr are equal, it is because from 55Cs to 87Fr only one shell increases but nuclear charge (No. of proton) increases by +32. (d) So effect of nuclear charge balanced the effect of increase in number of shell. Electronegativity of F > Cl Electron affinity of Cl > F but So Fluorine is called Black sheep element. (e) In group of IIB elements (Zn, Cd, Hg) value of electronegativity increases down the group, because of lanthanide contraction (f) In IIIA group, value of electronegativity increases down the group, because of transition contraction Electronegativity of Ga > Electronegativity of Al 24 Classification of Elements and Periodicity in Properties Note : HF has minimum bond length because of much difference in the electronegativities of H and F. (ix) Application of electronegativity : (A) Metallic and non metallic nature : Low electronegativity  Metals High electronegativity  Non metals Metallic character increases down the group but decreases along a period. (C) Bond energy : By increasing difference in electronegativity of bonded atoms, bond length decreases and hence bond energy increases (B) Bond length :    HF > HCl > HBr > HI Here EN = difference in electronegativities of bonded atoms Ex. HF < HCl < HBr < HI (D) Acidic strength of hydrides :  Order of stability of hydrohalides HF > HCl > HBr > HI In VA group – NH3 and order of acidic strength HF < HCl < HBr < HI PH3 AsH3  Thermal stability decreases Acidic character increases (E) Reactivity :   (b)      Ex. Note : HI is most reactive hydrohalides or strongest acid among all hydrohalides. Ex.  (F) Nature of bonds : (a) According to Hanny & Smith formula Ionic % = 16 (XA – XB) + 3.5 (XA – XB)2 Here XA = Electronegativity of A XB = Electronegativity of B Ex. If XA – XB > 2.1 Ionic % > 50% i.e. Ionic bond If XA – XB < 2.1 Ionic % < 50% i.e. covalent bond According to Gallis XA – XB > 1.7 Ionic XA – XB < 1.7 Covalent If XA = XB ; then A – B will be non polar. H—H, F—F If XA > XB and difference of electronegativities is small then A —— B bond will be polar covalent H2O ( H —— O —— H) If X A >> X B and X A – X B difference of electronegativities is high then A– ––— B+ bond will be polar or ionic Na+Cl– In HF, XA – XB = 1.9, which is more than 1.7, even then it is covalent compound. 25 Classification of Elements and Periodicity in Properties (G) Nature of hydroxides : Note: (a) As per Gallis, In AOH if electronegativity of A is BeO, Al2O3, ZnO, SnO, PbO, SnO2, PbO2, Sb2O3 more than 1.7 (Non metal) then it is acidic in nature. etc. are amphoteric oxides. (b) If electronegativity of 'A' is less than 1.7 (metal) then CO, H2O, NO, N2O etc. are neutral oxides. AOH will be basic in nature Ex. NaOH ClOH  XA 0.9 3.0 B2O 3 CO 2 N2 O 5 Nature Basic Acidic   (c) If XA – X0 > X0 – XH then AO bond will EN increase, acidic nature increasse. be more polar and   will break up as   A —— OH  A+ + OH– It shows basic   nature   (d) If XA – XO < XO – XH A — O ——H  H+ + AO– It shows Acidic nature Ex. In NaOH XO–XNa (2.6) > XO–XH (1.4) So hydroxide is basic In ClOH Acidic properties increases with increasing oxidation XO–XCl (0.5) < XO–XH (1.4) So hydroxide is state of an element acidic HClO4 > HClO3 > HClO2 > HClO (H) Nature of oxides : Consider an oxide AO H2SO4 > H2SO3 If XA – XO > 2.3 Basic oxide SO3 > SO2 If XA – XO = 2.3 Amphoteric oxide If XA – XO < 2.3 Acidic oxide HNO3 > HNO2 (a) Along a period acidic nature increases. N2O5 > N2O3 (b)Down the group basic nature increases Sb2O5 > Sb2O3 Li Be B C N O F (I) Hydrolysis of AX : Na Mg Al Si P S Cl Where A = Other element and X = Halogen  (a) If electronegativity of X > Electronegativity of A then XA – XO > 2.3 XA – XO = 2.3 XA – XO < 2.3 on hydrolysis product will be HX. Basic Amphoteric Acidic Ex. (BCl3), Electronegativity of Cl >Electronegativity of B ie. when in periodic table the distance between the element and oxygen increases, basic character increases. NO2 > ZnO > K2O PCl3 + 3HOH  3HCl + H3PO3  PCl3. Cl2 + H2O  2HCl + POCl3 acidic character decreases POCl3 + 3HOH  3HCl + H3PO4 26 Classification of Elements and Periodicity in Properties (b) If electronegativity of X < electronegativity of A then (B) Mulliken scale : According to Mulliken electronegativity is average value of ionisation on hydrolysis product will be HOX (hypohalous acid) potential and electron affinity of an element, Ex. Cl2O Here electronegativity of O > Cl   So On hydrolysis – (J) Nomenclature of inorganic compounds : Prefix —— less electronegative element Suffix —— More electronegative element Ex. Cl2O(Right) OCl2(Wrong) Relation between Pauling scale & Mulliken scale Xp = 0.336 (Xm – 0.615) Xp = Electronegativity given by pauling Xm = Electronegativity given by mulliken  If ionisation potential and electron affinity are given in eV, then electronegativity by Mulliken on Pauling scale will be In Dichloroxide the electronegativity of Cl is less than 'O' i.e. why Cl is in prefix position. OF 2 Oxygen difluoride ICl Iodine chloride    If ionisation potential and electron affinity are given in K.cal/mole then   Exception (Here H is less electronegative  but it is suffix) (C) Allred Roschov's scale : As per Roschov, (x) Electronegativity Scale : electronegativity is the force by which nucleus of an atom attracts electron which are on the covalent (A) Pauling Scale (Bond energy scale) : radius. NH3    Here AB is resonance energy of AB molecule This equation gives difference in electronegativity values.  or 27