CHM 101 Introductory Inorganic Chemistry March 2021 PDF

COURSE GUIDE CHM 101 INTRODUCTORY INORGANIC CHEMISRTY Course Team Professor John Kanayochukwu Nduka (Course Reviewer) – NAU Awka NATIONAL OPEN UNIVERSITY OF NIGERIA CHM 101 COURSE GUIDE © 2021 by NOUN Press National Open Un...

COURSE GUIDE CHM 101 INTRODUCTORY INORGANIC CHEMISRTY Course Team Professor John Kanayochukwu Nduka (Course Reviewer) – NAU Awka NATIONAL OPEN UNIVERSITY OF NIGERIA CHM 101 COURSE GUIDE © 2021 by NOUN Press National Open University of Nigeria Headquarters University Village 91, Cadastral Zone Nnamdi Azikiwe Expressway Jabi, Abuja Lagos Office 14/16 Ahmadu Bello Way Victoria Island, Lagos e-mail: [email protected] URL: www.nou.edu.ng All rights reserved. No part of this book may be reproduced, in any form or by any means, without permission in writing from the publisher. Printed by NOUN Press Published 2021 ISBN: 978-978-058-076-6 ii CHM 101 COURSE GUIDE CONTENTS PAGE Introduction …………………………………….. iv What you will learn in this course………………. iv Course Aims……………………………………... iv Course Objectives……………………………….. v Working through this course…………………….. v Course Materials ………………………………... v Study Units……………………………………… v Textbooks……………………………………….. vii Assessment……………………………………… vii Summary………………………………………... vii iii CHM 101 COURSE GUIDE INTRODUCTION Inorganic Chemistry is the study of chemical elements and their compounds excluding compounds of carbon; but compounds of carbon and oxygen, Sulphur and metals are generally included in inorganic chemistry. An element is a substance which cannot be resolved into two or more simpler substances following any known chemical process. Combination 0f two or more elements chemically form a new substance called compounds. Examples, water (formed from hydrogen and oxygen), common salt (formed from sodium and chlorine), carbon dioxide (from carbon and oxygen). Since scientist systematize the knowledge gained through observation and experiments. Development of periodic table and periodic law is an example of such attempt that has brought order in studying an expansive branch of chemistry covering over a hundred element. It is therefore pertinent to start the study of inorganic chemistry with the focus on the periodic table. CHM 101 – Introductory inorganic chemistry is a two (2) credit hour course of seventeen (17) units. The course is designed to equip the student with in- depth knowledge of the periodic classification of element, properties of element according to groups and periods. The course is designed to deal with periodic table and periodic law, electronic configuration, atomic radii, ionization energy, affinity of electrons, electronegativity, hydrogen, its reaction and compounds, Alkali and Alkaline earth metals compounds, physical and chemical properties etc. This course guide gives the student a brief overview of the course content, course duration, and course materials. WHAT YOU WILL LEARN IN THIS COURSE The student will learn about the periodic table, how it came to be, and the periodic law. Scientist that contributed to the development of periodic table. Hydrogen, chemical properties and reactions. Alkali and Alkaline earth metals, their physical and chemical properties. Reactions and uses. COURSE AIMS The aim of this course is tailored towards teaching the students to understand periodic classification of elements, various contributors to the development of periodic table. Modern periodic law, electronic configuration and rules governing its atomic radii, ionization energy, affinity of electrons, electronegativity, Hydrogen, its properties, its bonding, manufacture, ionic iv CHM 101 COURSE GUIDE salt, Alkali and Alkaline earth metals, their general physical and chemical properties, compounds and reactions. COURSE OBJECTIVES To make sure that this course achieves its aim, the stratified course objectives is shown below. The entire objectives are specific and peculiar to the course objectives. The course objectives are stated below: i. To understand the scope and the basic principle of introductory inorganic chemistry ii. To understand the elemental arrangement in the periodic table, electronic configuration of element and periodic law. iii. To understand the work of scientist that led to the development of the periodic table, properties and reactions of elements according to positions in the periodic table, atomic radius. iv. To understand hydrogen, its manufacture, reactions, its ionic salts-like hydrides, bonding, alkali and alkaline earth metals, their compound and reactivity. WORKING THROUGH THIS COURSE This contains some packages that will be given to the students at the beginning of the semester. It includes the course material. You will be required to participate fully in the continuous assessment and the final written examination in three areas such as TMA, MCQ and FBQ etc. The seventeen units of the course packaged for you in modules are summarized below COURSE MATERIALS 1. Course Guide 2. Study Units v CHM 101 COURSE GUIDE STUDY UNITS The following are the three modules and the seventeen units contained in this course. Module 1 Unit 1 Periodic Table Unit 2 Modern Periodic Law Unit 3 Electronic Configuration Unit 4 Atomic Radius Unit 5 Ionization Energy Unit 6 Electron Affinity Unit 7 Electronegativity Module 2 Unit 1 Hydrogen Unit 2 Manufacture of Hydrogen Unit 3 Ion and Salt-like Hydride Unit 4 Hydrogen Bonding Module 3 Unit 1 General Physical and Chemical Characteristics of the Alkaline Metals Unit 2 Compounds of Alkali Metals Unit 3 Solvation of Alkaline Metal Ions Unit 4 Alkaline Earth Metals Unit 5 Reactivity of Alkaline Earth Metals Unit 6 Complexing Behaviour of Alkaline Earth Metals In module 1, introduction to periodic classification, attempt and contribution of J.W Dobereiner, Ade Chancourtois, John Newlands etc. Unit 2 discussed modern periodic law and nomenclature of elements having Z > 100; Unit 3 deals with the rules of filling orbitals, configuration of elements, ion and division of elements into blocks. Atomic radii, covalent radius, periodicity, periodicity in ionization energy etc. were discussed in unit 4 and 5 while unit 6 and 7 discussed factors affecting electron affinity, Pauling electronegativity, Maulliken-Jafffe electronegativity and Alfred Rochow electronegativity. vi CHM 101 COURSE GUIDE Module 2 contains units 8, 9, 10 and 11 and contains extensive discussion of hydrogen, its position in the periodic table, its isotopes, manufacture, properties, uses, ionic or salt-like hydrides and metallic hydride, Hydrogen bonding, its effect on boiling and melting point, water solubility and polarizing power of hydrogen ion. Unit 12 of module 3 took a look at the general physical and chemical properties of alkali and alkaline earth metal, their uses. Unit 13 discussed the compounds of alkali metals, oxides of hydrogen. Unit 14 explained solvation of alkaline metal ions, their complexation behavior and anomalous nature of lithium. Unit 15 and 16 dwelt on alkaline earth metals and their properties reactivity, occurrence, extraction, lattice energy, hydration energy and thermal stability of oxysalts. The course ends with unit 17 of module 3 by explaining complexation behaviour of alkaline earth metals and anomalous nature of Beryllium. TEXTBOOKS AND REFERENCES There are several books and other materials that treated Inorganic Chemistry 101. A good number of them are written down in the references. The internet provides a lot of useful information concerning the course, therefore, the student is encouraged to use the internet and NOUN e-library were possible. ASSESSMENT There are two aspects of assessment for this course; the tutor-marked assignment (TMA) and end of course examination. TMAs shall constitute the continuous assessment component of the course. They will be marked by the tutor and shall account for 30% of the total course score. Each learner shall be examined for four TMAs before the end of course examination. The end of course examination shall constitute 70% of the total course score. SUMMARY This course is necessary as an introductory Inorganic Chemistry course. It opens up and introduces the student into the wider knowledge of the Chemistry of all the elements as classified in the periodic table, the historical development of the periodic table, the special place of Mendeleev in the development of modern periodic table and periodic law. Chemical and vii CHM 101 COURSE GUIDE physical properties of elements as influenced by their position in the periodic table (Groups and Periods), their compounds, extraction, manufacture and uses. We wish you an outstanding success in this course and others as you climb the ladder of knowledge. We hope that as you study hard to pass your examinations excellently, you will also allow the knowledge you acquire to have a positive influence on the society and environment. viii  MAIN COURSE CONTENTS PAGE Module 1…………………………………………………… 1 Unit 1 Periodic Table ………………………………. 1 Unit 2 Modern Periodic Law………………………... 11 Unit 3 Electronic Configuration…………………….. 22 Unit 4 Atomic Radius……………………………….. 35 Unit 5 Ionization Energy…………………………….. 52 Unit 6 Electron Affinity……………………………… 62 Unit 7 Electronegativity……………………………… 67 Module 2……………………………………………………. 75 Unit 1 Hydrogen ……………………………………. 75 Unit 2 Manufacture of Hydrogen…………………… 83 Unit 3 Ion and Salt-like Hydride……………………. 93 Unit 4 Hydrogen Bonding…………………………… 98 Module 3……………………………………………………. 106 Unit 1 General Physical and Chemical Characteristics of the Alkaline Metals………………………… 106 Unit 2 Compounds of Alkali Metals…………………. 118 Unit 3 Solvation of Alkaline Metal Ions…………….. 126 Unit 4 Alkaline Earth Metals………………………… 133 Unit 5 Reactivity of Alkaline Earth Metals………….. 141 Unit 6 Complexing Behaviour of Alkaline Earth Metals………………………………………… 149 CHM 101 MODULE 1 MODULE 1 Unit 1 Periodic Table Unit 2 Modern Periodic Law Unit 3 Electronic Configuration Unit 4 Atomic Radius Unit 5 Ionization Energy Unit 6 Electron Affinity Unit 7 Electronegativity UNIT 1 THE PERIODIC TABLE 1.0 Introduction 2.0 Intended Learning Outcomes (ILOs) 3.0 Main content 3.1 Beginning of Classification 3.2 Attempt made by J.W Dobereiner 3.3 Attempt made by Ade Chancourtois 3.4 Attempt made by John Newlands 3.5 The work of Lothar Meyer 3.6 Mendeleevs Periodic law 4.0 Conclusion 5.0 Summary 6.0 Tutor Mark Assignment 7.0 References/Further Readings 1.0 INTRODUCTION You are aware that scientists, from the very beginning have attempted to systematize the knowledge they gain through their observations and experiments.. Development of the periodic law and the periodic table of the elements is one of such attempt. This has brought order in the study of the vast chemistry of more than a hundred elements known now. It is therefore quite natural that you should begin your study of inorganic chemistry with the study of the periodic table. In this unit, you will be starting from the very beginning, that is, with the very first attempt made at classification of the elements. As early as 1815 Prout advanced the hypothesis that all elements were formed by the coalescence of hydrogen atoms, on inaccurate evidence that all atomic masses were whole numbers but it was Berzelius who showed that ‘atomic mass’ of chlorine was not 35 nor 36 but 35.5 far from a whole number. The discovery of isotope proved that chlorine was a mixture of containing two different kinds of atoms of mass 35 and 37. 1 CHM 101 INTRODUCTION INORGANIC CHEMISTRY By the mid-nineteenth century, more than 60 elements were known and many m o r e were being discovered. The rate of discovery of the new elements was so fast that the chemists started wondering "where it would all lead to". Has nature provided a limit to the number of elements? And if so, how would one know about it? During this period, it was also realized that certain groups of elements exhibited similar physical and chemical properties. Was it a mere coincidence or did a relationship exist among theproperties of the elements? Attempts reply such probing questions ultimately resulted in the development of the periodic table. 2.0 OBJECTIVES By the end of this unit, you should be able to:  List accurately at least two scientists who attempted to classify the elements into periods.  Write with at least 70% accuracy, brief accounts of the attempts made by the two scientists and the result of these attempts.  State Mendeleev's periodic law.  State the property used by Mendeleev to classify the elements in his periodic table.  Demonstrate an understanding of Mendeleev's law by applying it to predict properties of undiscovered elements. 3.0 MAIN CONTENT 3.1 The Beginning of Classification One of the earliest attempts to classify elements was to divide them into metals and non-metals. Metallic elements we all know have certain properties which include 1. Having lustrous shinning appearance, such as iron 2. Malleability, meaning they can be beaten into thin sheets such as is done when buckets are being produced 3. Metallic elements can so be drawn into wire such as is done when making electric wire. This property is known as ductility. 4. They can also conduct heat and electricity. If you hold a piece of metal in your hand and put on end into fire or in contact with any hot object, your hand will feel the heat as it travels from the point of contact with heat through the metal to your hand. Similarly, if you hold one end of the metal through which an electric current is passed, you will be jolted by the current which travel through the metal to your hand. 5. Metallic elements also form basic oxides 2 CHM 101 MODULE 1 In contrast to metallic elements, non-metallic elements have no characteristic appearance. They are brittle that is they break easily. They are poor conductors of electricity and heat. They form acidic oxides. As more elements were discovered and knowledge of physical and chemical properties were refined, it became clear that within these two divisions of elements, there existed families of elements whose properties varied systematically from each other. Furthermore, certain elements, the metalloids possessed properties intermediate between the two divisions. Thereore were made to search for other classifications. 3.2 Attempts Made by J W Dobereiner In 1829 J W Dobereiner observed that there exist certain groups of three elements which the called TRIADS. He also observed that elements in triad not only had similar properties, but also the atomic weight of the middle element was approximately an average of the atomic weights of the other two elements of the triad. A few examples cited by him were: Li, Na, K, Ca, Sr, Ba, S, Se, Te and Cl. Br, I, Al though, Doberieiner's relationship seems to work only fora few elements, He was the first to point out a systematic relationship among the elements. This was followed by Cannizzaro’s unambiguous atomic mass with the consequent allotment of valencies to atoms as a measure of combining power. In-text Question State Prout’s (1815) hypothesis of the elements explain how the error was corrected. 3.3 Attempts Made by A. Dechancourtois In 1862, A. DeChanourtois arranged the elements that were known at that time in order of increasing atomic weight on a line which spiralled around a cylinder from bottom to top. 3.4 Attempts Made by John Newlands In 1864, John Newlands, an English Chemist reported his "LAW OF OCTAVES" He suggested that if the elements were arranged in order of increasing atomic weight, every eighth element would have properties similar to the first element. It brought the lithium-sodium- potassium triad together but failed to allow bigger octave if dealing with 3 CHM 101 INTRODUCTION INORGANIC CHEMISTRY heavier elements. For example, He arranged the elements in the following manner. Table 1: John Newland’s arrangement of the element Element Li Be B c N 0 F AtWt 7 9 11 12 14 16 19 Element Na Mg Al Si P S Cl AtWt 23 24 27 29 31 32 35.5 Element K Ca Ti Cr AtWt 39 40 48 32 Thus we see K resembles Na and Li, Ca resembles Mg and Be, Al resembles B, Si resembles C and so on. He called it the "Law of octaves" because he says the cycle of repetition shown by the elements is like that shown by octaves of music where every eight note resembles the first in octaves of music. Newlands "Law of octaves" was rejected for two reasons. Firstly, it did not hold good for elements heavier than Ca. Secondly, he believed that there existed some mystical connection between music and chemistry. 3.5 The Work of Lothar Meyer Lothar Meyer and Dmitri Mendeleev whom you will read about next played key role in the development of the periodic law as it is known today. In 1869, Lothar Meyer reported that when physical properties like atomic volume, boiling point etc. were plotted against atomic weight, a periodically repeating curve was obtained in each case. Figure 1 is a graph showing the variation in atomic volume with atomic number. Lothar Meyer calculated the atomic volumes of the known elements, that is the volume in cm3 occupied by 1mole of the elements in the solid state Atomic volume = Mass of one Mol ̷ Density (Lothar Meyer also obtained semi curve by plotting atomic volume versus atomic weight) The atomic volume behaviour is p e r i o d i c. It goes through circles, dropping from a sharp maximum to a minimum and then sharply rising again. Each of the cycles is called a period. The location of element on the peak or in the troughs has an important correlation with their chemical 4 CHM 101 MODULE 1 reactivity. The elements of the peaks (example alkali metals) are the most reactive. Those in the troughs (example noble metals) are characteristically less reactive. Figure 1. Periodic dependenceof atomic volume on atomic number 3.6. Mendeleev's Periodic Law In contrast to Lothar Meyer, Mendeleev used chemical properties likevalence and formulae of hydrides, chloride, and oxides of the elements to illustrate his periodic law. According to Mendeleev's periodic law, if the elements are arranged sequentially in the order of increasing atomic weight, a periodic repetition, that is, periodicity in properties is observed. Mendeleev arranged elements in horizontal rows and vertical columns in order of increasing atomic weight so that the elements having similar properties were kept in the same vertical column. 5 CHM 101 INTRODUCTION INORGANIC CHEMISTRY Table 2. Mendeleev periodic table of against each element is the value of atomic weight Ser Gro Gro Grou Grou Gro Grou Grou ies up I up II p III p IV up V p VI p VII Group VIII - - - RH4 RH3 RH2 RH - R2O R1O RO R2O3 RO2 IO2 R1O1 RO4 6 1 H=1 Li= Bc= B=1 C=1 N=1 O=1 2 7 9.4 1 2 4 6 F=19 Na= Mg Al=2 Si=2 P=3 Cl=3 3 23 =24 7.3 8 1 S=32 5.5 K=3 Ca= Sc=4 Ti=4 V=5 Cr=5 Mn= Fe=56, Co=59 4 9 40 4 8 1 2 55 Ni=59 Cu= Zn= Ga= Ge= As= Sc=7 Bi=8 5 63 65 68 72 75 8 0 Ru=104, Rb= Sr= Yt=8 Zr=9 Nb= Mo= Fm= Rh=104 6 85 87 8 0 94 96 100 Pd=106, Ag= Cd= In=1 Sn=1 Sb= Te=1 I=12 7 108 112 13 18 122 25 7 Cs= Ba= Da= Ce= -- 8 133 137 138 148 - - - 9 - - - - - - - -- Os=195, Er=1 La=1 Ta= W=1 Ls=195, 10 - - 78 80 182 84 Pi=198 Au= Hg= Ti=2 Pb=2 Bi= 11 199 200 04 07 208 Th= U=2 -- 12 - - 231 - 40 - Though Newlands and Lothar Meyer also contributed in developing the periodic laws, the main credit goes to Mendeleev because of the following reasons: a. He included along with his table, a detailed analysis of the properties of all known elements and correlated a broad range of physical and chemical properties with atomic weights. 6 CHM 101 MODULE 1 b. He kept his primary goal of arranging similar elements in the same group quite clear. Therefore, he was bold enough in reversing the order of certaine elements. For example, iodine with lower atomic weight than that of tellurium (group VI) was placed in group VII along with fluorine, chlorine and bromine because of similarities in properties. c. He also corrected the atomic weight of certain elements to include the min proper groups. For example, he corrected the atomic weight of beryllium (from 13.5 to 9) and indium (from 76 to 114) without doing any actual measurement. His competence was proved correct as Be and La with equivalent weight of 4.5 and 38 respectively are actually bivalent and trivalent. d. Keeping to his primary goal of arranging similar elements in the same vertical column (group), he realized that some of the elements were still undiscovered and therefore left their places vacant in the table and predicted their properties. He predicted the existence in nature of over ten new elements and predicted properties of three of them, example eka-boron (scandium), eka aluminium (gallium) and eka silicon (germanium) from the properties of known elements surrounding them. When these elements were eventually discovered, Mendeleev prediction proved to be amazingly accurate. This you can see for yourself by comparing the prediction and observed properties of eka- alurninium (gallium) and eka silicon (germanium) given in table 2. The validity of Mendeleev periodic law was dramatically and conclusively proven by the discovery of three out of the more than ten elements predicted by Mendeleev. The first to be discovered was eka-aluminiurn which was discovered by Lecoq de Boisbaudran in 1875. 7 CHM 101 INTRODUCTION INORGANIC CHEMISTRY Table 3. Comparison of predicted and observed properties of eka Aluminum (gallium) and eka Silicon (germanium) Property Predicted Observed Predicted Observed for by for by germanium Mendeleev gallium Mendeleev for eka- for aluminum ekasilicon Atweight 68 69.72 72 72.59 Density 6.0 x l03 5.9 x103 5.5 x l03 5.3 x103 -3 (kgm ) Melting Low 302.8 High 1220k point/k Reaction Slow Slow Slow Reacts with with acids concentrated acids &alkalines & alkaline Formula of E2O3 Ga2O3 - - oxide Density of 5.5 x103 5.8 xl03 - - oxide (kgm-3) Formula ECl3 GaCl3 EsCl4 GeCl4 for chloride Boiling Volatile 474 373 357 point of chloride (k) Lecoq de Boisbaudran called the element gallium and said its density was 4.7 x 103 kgm-3. Mendeleev on hearing this wrote to Lecoq de Boisbaudran telling him that everything he said about the new element was correct except its density. On further position of the metal, lecoq de Biosbandran discovered that Mendeleev was right that the density of gallium was 5.8xlO3 kg just like it had been predicted by Mendeleev. (Table 3) Further proof of the law came via the works of Lars Fredrick Nilson who discovered Scandium and Winkler who discovered germanium. Both elements were found to have properties corresponding to those of earlier predicted for them by Mendeleev. 8 CHM 101 MODULE 1 The development of the periodic law is an excellent example where careful observation, critic analys of available data without any pre- conceived notions and scientific foresight led to the discovery of a fundamental law of nature. Thus when Mendeleev arranged elements in order of increasing atomic weights, he critically analyzed the properties of the then known elements. He discovered that the properties of any element are an average of the properties of its neighbours in the periodic table. On this basis, he predicted the properties of undiscovered elements representing the gaps in the table. The many gaps in the table which Mendeleeff correctly predicted would eventually be filled by the discovery of more elements. No gaps remain to be filled, therefore the properties of many of these element are in close agreement with those predicted by Mendeleeff. In modern periodic table, the ordering of the element is based on atomic numbers rather than atomic masses. This accommodated the discovery of the noble gasses, 14 rare earths (the first inner transition series called the lanthanides) and 11 man-made elements (part of the second inner transition series called the actinides) In-text Question Question State the point that makes Mendeleev’s periodic law stronger than Lother Meyer’s SELF-ASSESSMENT EXERCISES i. Give the names of the scientist whose work contributed to the development of the periodic table? ii. The law of octave was developed by_____ 4.0 CONCLUSION In conclusion, the Scientists tradition of recording and system knowledge gained through observations and experiments has enabled us to learn about the fundamental laws governing the arrangement of elements. Mendeleev prediction proved to be amazingly accurate. This you can see for yourself by comparing the prediction and observed properties of eka- alurninium (gallium) and eka silicon (germanium). 5.0 SUMMARY In summary, we have learned the following in this unit, these are: 1. Scientists have always tried to systemize the knowledge they gain. 9 CHM 101 INTRODUCTION INORGANIC CHEMISTRY 2. That the effort to reveal the secrets of the periodic table were led by the scientist J W Dobreiner, A de chanourtois, John Newlands, Lothar Meyer and Dmitir Mendeleev. 3. That the works of Dmitir Mendeleev formed the basis of the modern Periodic law. 4. That according to Mendeleev, the properties o f any element a r e an average of the properties of its neighbours in the periodic table. 6.0 TUTOR MARK ASSSIGNMENT 1. (a) What property did Mendeleev use to classify the element in his periodic table? (b) Enumerate four defects in the Mendeleev's periodic table 2. Assuming that the element Ca had not been discovered, predict using the properties of the Known element surrounding Ca its own properties such as its atomic weight and density. 7.0 REFERENCES/ FURTHER READING J. G Wilson, A. B. Newell (1971) General and Inorganic Chemistry (2nd edition) Published by Cambridge University Pres F. A Cotton, G. Wilkinson and P. L. Gayus (1995) Basic Inorganic Chemistry. (3nd edition) John Wiley and Sons Published Gary L. Miessler, Paul J. Fischer and Donald A. Tour (2014) Inorganic Chemistry. 5th edition. Pearson Publishers. ISBN 10:0-321- 81105-4 J. D Lee (2016) Concise Inorganic Chemistry. (5th edition) John Wiley Publishers 10 CHM 101 MODULE 1 UNIT 2 MODERN PERIODIC LAW 1.0 Introduction 2.0 Objectives 3.0 Main content 3.1 Modern Periodic Law 3.2 The long form of The Periodic Table 3.3 Nomenclature of Element having Z >100 4.0 Conclusion 5.0 Summary 6.0 Tutor Mark Assignment 7.0 References/Further Readings 1.0 INTRODUCTION In the first unit, you learned about efforts made by several scientists to systematize the knowledge they gained through their observations and experiments. The effort of these scientists resulted in the formation of the periodic table and the periodic law. You learned about the works of such great pioneers as J. N Dobereiner, A de Chancoutois, John Newlands, Lothar Meyer and Dmitir Mendeleev. The periodic table of today has many similarities with that formed by Mendeleev, but differs from the Mendeleev table in some significant ways which we shall see in the next unit. Also in the past, element were named by their discoverers. In some cases, such a practice has led to disputes between scientists who have discovered the same elements working independently in different parts of the world. This has prompted the international union of pure and applied chemists’ to device a method for naming newly discovered elements 2.0 OBJECTIVES By the end of this unit, you should be able to:  State the modem periodic law.  Explain the relative positions of K and Ar, Co and Ni and Te and I on the periodic table.  State the relationship between the atomic number and the Periodic classification of elements  Apply IUPAC nomenclature rules in naming new elements having Z>100 11 CHM 101 INTRODUCTION INORGANIC CHEMISTRY 3.0 MAIN CONTENT 3.1 MODERN PERIODIC LAW In the previous section, you studied how D. Mendeleev classified elements and formed his periodic table. You must have noticed that there were anomalies in Mendeleev's original periodic table. There was for example no place for lanthanides and actinides and in some instances, elements of higher atomic weight were placed before those of lower atomic weights example Co before Ni and Te before I. He could not predict the existence of noble gases, nor could he properly place hydrogen. Between 1869-1907Mendeleev tried to improve his table. However, the most significant improvement of his periodic table came through the discovery of the concept of atomic number in 1913 by Henry Moseley, who suggested that the atomic number of an element is a more fundamental property than its atomic weight. Mendeleev's periodic law was therefore accordingly modified. This is now known as the MODERN PERIODIC LAW and can be stated as "the properties of elements are periodic functions of their atomic numbers" Arrangement of the elements in order of their increasing atomic number removes most of the anomalies of Mendeleev's periodic table. The positions of K and Ar, Co and Ni, Te and I do not remain anomalous any longer since atomic number not weight is used in arranging the elements. As isotopes of an element have the same atomic number, they can all be placed at one and the same place in the periodic table. We know that the atomic number cannot be fractional. It increases by the integer from one element to the next. It has thus placed a limit on the number of elements. Today, 109 elements (from 1 to 109) have been discovered. And any more elements that may be discovered in future will be beyond 109. Table 4. Shows the modem periodic table in the form devised by Mendeleev 12 CHM 101 MODULE 1 Periods A I B A VII B A VIII B 1 1 1 H H 2 He 1.008 A II B A III B A IV B AVB A VI B 1.008 4.003 3 4 5 6 7 8 10 2 Li Be B C N O 9 F Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 18 3 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl Ar 22.99 24.31 26.98 28.09 30.97 32.06 35.45 39.95 22 26 27 28 19 K 20 Ca 21 Sc Ti 23 V 24 Cr 25 Mn Fe Co Ni 39.10 40.08 44.96 47.80 50.94 52.00 54.94 55.85 58.71 58.93 4 35 36 29 Cu 30 Zn 31 Ga 32 33 As 34 Se Br Kr 63.54 65.37 69.72 Ge 72.59 74.92 78.96 79.91 83.80 44 45 46 37 Rb 38 Sr 39 Y 40 41 Nb 42 Mo 43 Tc Ru Rh Pd 85.74 87.62 88.91 Zr 91.22 92.91 95.94 98.91 101.07 102.91 106.4 5 54 47 Ag 48 Cd 49 In 50 51 Sb 52 Te 53 I Xe 107.87 112.40 114.82 Sn 118.96 121.75 127.60 126.90 131.30 13 CHM 101 INTRODUCTION INORGANIC CHEMISTRY 72 76 77 78 Os Ir 55 Cs 56 Ba 57 La* Hf 73 Ta 74 W 75 Re Pt 190.2 192.2 132.91 137.54 138.91 178.49 180.95 183.85 186.2 195.09 6 79 Au 80 Hg 81 Ti 82 83 Bi 84 Po 85 At 86 197.97 200.59 204.37 Pb 207.19 208.98 120 120 Rn 222 104 Unq 105 Unp 106 Unh 107 Uns 108 109 Uno 7 87 Fr 88 Ra 89 Ac** Une 223 226.03 227.03 *Lanthanides 58 59 60 61 62 63 64 65 66 67 68 69 70 71 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 140.12 140.91 144.24 146.92 150.35 151.96 157.25 158.92 162.50 164.93 167.26 168.93 173.04 174.97 **Actinides 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr 232.04 231.04 238.03 237.05 239.05 241.06 247.07 249.08 251.08 254.09 257.10 258.10 255.0 257.0 14 CHM 101 MODULE 1 3.2 Long Form of the Periodic Table You have now seen that in the modern form of Mendeleev periodic table, elements are arranged in seven horizontal rows and eight vertical columns. Normal and transition elements belonging to A and B sub group of a group were placed in one and the same column of the table. For example, Sc and Ga both in group IIIA and B and Ti and Ge both in group IVA and B. In the long form of the periodic table (see Table 4) elements are arranged in eighteen vertical columns by keeping the Elements belonging to A and B sub groups in separate columns. Note that in the new arrangement, Sc is in group IIIB whereas Ga is now placed in group IIIA. You would have also noticed that the groupVIIIB of Mendeleev's periodic table contains three triads Fe, Co, Ni, (4th period), Ru, Rh, Pd (5th period) and O, S, Ir, Pr (1st period). In the long form of the table, each element of the triad is kept in a separate column. So the group VIIIB occupies three columns of the table. You can see therefore that, the long form of periodic table is an extension of the modern periodic table 15 CHM 101 INTRODUCTION INORGANIC CHEMISTRY Table 5: Long form of the periodic table 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 IA IIA IIIB IVB VB VI VII VIII IB IIB IIIA IVA VA VIA VIIA VIII B A 100 https://www.youtube.com/watch?v=8MxXNcKtCII 21 CHM 101 INTRODUCTION INORGANIC CHEMISTRY UNIT 3 ELECTRONIC CONFIGURATION 1.0 Introduction 2.0 Objectives 3.0 Main Content 3.1 Rules governing the filling of electrons in orbitals 3.2 Electronic configuration of all the elements in the periodic table 3.3 Electronic configuration of ions 3.4 Electronic configuration and division of Element into blocks 4.0 Conclusion 5.0 Summary 6.0 Tutor Mark Assignments 7.0 References/Further Readings 1.0 INTRODUCTION In the last unit, you learned about the periodic table. The periodic law gives rise to the periodic arrangement of the element according to their atomic numbers, which indeed means according to the number of electrons in their orbital. In the next unit you will be studying the distributions of the electrons within the atom (the electronic configuration), and how they govern the properties of the elements. The electronic configurations of isolated atoms of elements are usually verified experimentally by a detailed analysis of atomic spectra. In this unit, we are not discussing these methods. Instead we are going to discuss the process of filling in of electrons in the various atomic orbitals. 2.0 OBJECTIVES  State the principle involved in determining which electron goes into which atomic orbital  Fill out correctly electrons in a given atom once given the number.  List the four blocks of elements on the periodic table and determine to which block an element  belongs if given the electronic configuration. 3.0 MAIN CONTENT 3.1 Electronic Configuration of Atoms Rules governing the filling of electrons in orbital 22 CHM 101 MODULE 1 The electronic configuration of atoms can be predicted with the help of AUFBAU or the building up process. In the aufbau process, it is assumed that there exists asset of empty hydrogen like orbital around the nucleus of an atom. The electronic configuration of the atom in the ground state is then derived by adding electrons one at a time to the orbitals of the lowest energy in the sequence shown by arrows in figure 2 Fig 2: Order or filling of atomic orbitals in poly electronic atoms The order in which the orbitals are filled as shown in figure 2, is governed by the n+1 rule. According to this rule, in the building up of electronic configuration of the elements, the sub-shell with the value of n+l fills first. This rule reminds us that the energy of sub shells of multi electron atoms 23 CHM 101 INTRODUCTION INORGANIC CHEMISTRY depends upon the value of both the quantum numbers n and 1, but mainly on the value of n. For example, to determine whichof the sub shells 5s or4p fills first. We follow the rule thus: for the 5s sub shell the value n+1 =5 +0=5; for 4p sub shell also the value of n+1 =4+ 1 =5, but the 4p sub shell has the lower. Value of the principal quantum number n and therefore it fills first. Filling of electrons in orbitalsis also governed by Paulis Exclusion Principle and Hund's Rule. The Pauli Exclusion Principle states that no two electrons in the same atom can have the same values for all the four quantum numbers while hund’s rule states that electrons occupy each levels singly before electron pairing takes place. According to the Exclusion Principle, no two electrons in the same atom can have the same value of n, l and mi, they will differ in their m, values. In order words, an orbital can have at the most two electrons of opposite spin. Since there is only 1s -orbital for any given value of n, it can contain only two electrons. However, the three p orbitals for any given value of n can contain six electrons, the five d orbitals, for any given value of n can hold a total of ten electrons and the seven f orbitals can have fourteen electrons. Permitted combinations of all the four quantum numbers for the electrons in different orbitals are given below in table 7 Hund's Rule Table 7. Permitted combinations of quantum numbers for S.P, d and f orbitals M T m1 m1 Common Number of name Electron 1 0 0 ±1/2 1r 2 2 0 0 ±1/2 2r 2. 2 1 -1 ±1/2 2p 6 0 ±1/2 +1 ±1/2 3 0 0 ±1/2 3r 2 3 1 -1 ±1/2 3p 6 0 ±1/2 +1 ±1/2 3 2 -2 ±1/2 3d 10 -1 ±1/2 0 ±1/2 +1 ±1/2 +2 ±1/2 4 0 0 ±1/2 4s 2 4 1 -1 ±1/2 4p 6 24 CHM 101 MODULE 1 0 ±1/2 +1 ±1/2 4 2 -2 ±1/2 4d 10 -1 ±1/2 0 ±1/2 +1 ±1/2 +2 ±1/2 4 3 -3 ±1/2 4f 14 -2 ±1/2 -1 ±1/2 0 ±1/2 +1 ±1/2 +2 ±1/2 +3 ±1/2 5 0 Hunds rule of maximum multiplicity states that, as far as possible in a given atom in the state, electrons in the same sub shell will occupy different orbitals and will have parallel spins. That means that when electrons are added to orbitals of the same energy such as three p orbitals or five d orbitals, one electron will enter each of the available orbital, two electrons in separate orbitals feel less repulsion than two electrons paired in the same orbitals. For example, carbon in the ground State has the configuration 1S22S22Px12Py1 rather than 1S22S22Px2. so far you have studied the rules governing the filling of electrons in the orbitals of atoms. We shall now consider the electrons configurations of all the elements in the periodic table, these are given in table 8 Table 8. Ground state electronic configuration of gaseous atoms Configuration as Core plus outermost Z Symbol orbital 1 H |S1 2 He |S2 , or He | 3 Li [He]2S1 4 Be [He]2S2 5 B [He]2S22p1 6 C [He]2S22p2 7 N [He]2S22p3 8 O [He]2S22p4 9 F [He]2S22p5 10 Ne [He]2S22p6 or [Ne] 11 Na [Ne]3S1 12 Mg [Ne]3S2 25 CHM 101 INTRODUCTION INORGANIC CHEMISTRY 13 Al [Ne]3S23P1 14 Si [Ne]3S23P2 15 P [Ne]3S23P3 16 S [Ne]3S23P4 17 Cl [Ne]3S23P5 18 Ar [Ne]3S23P6 or [Ar] 19 K [Ar]4S1 20 Ca [Ar]4S2 21 Sc [Ar]3d14s1 22 Ti [Ar]3d24s2 23 V [Ar]3d34s2 24 Cr [Ar]3d44s2 25 Mn [Ar]3d54s1 26 Fe [Ar]3d64s2 27 Co [Ar]3d74s2 28 Ni [Ar]3d84s2 29 Cu [Ar]3d104s1 30 Zn [Ar]3d104s2 31 Ga [Ar]3d104s14p1 32 Ge [Ar]3d104s14p2 33 As [Ar]3d104s14p3 34 Se [Ar]3d104s14p4 35 Br [Ar]3d104s14p5 36 Kr [Ar]3d104s14p6 or Kr 37 Rb [Kr]5S1 38 Sr [Kr]5S2 39 Y [Kr] 4d15S2 40 Zr [Kr] 4d25S2 41 Nb [Kr] 4d45S1 42 Mo [Kr] 4d55S1 43 Tc [Kr] 4d65S2 44 Ru [Kr] 4d75S1 45 Rh [Kr] 4d85S1 46 Pd [Kr] 4d10 47 Ag [Kr] 4d105S1 48 Cd [Kr] 4d105S2 49 In [Kr] 4d105S25p1 50 Sn [Kr] 4d105S25p2 51 Sb [Kr] 4d105S25p3 52 Te [Kr] 4d105S25p4 53 I [Kr] 4d105S25p5 54 Xe [Kr] 4d105S25p6 or [Xe] 55 Cs [Xe]6S1 56 Ba [Xe]6S2 57 La [Xe] 5d16S2 26 CHM 101 MODULE 1 58 Ce [Xe] 4f15d26S2 59 Pr [Xe] 4f56S2 60 Nd [Xe] 4f56S2 61 Pm [Xe] 4f56S2 62 Sm [Xe] 4f56S2 63 Eu [Xe] 4f56S2 64 Gd [Xe] 4f15d16S2 65 Tb [Xe] 4f56S 66 Dy [Xe] 4f56S2 67 Ho [Xe] 4f516S2 68 Er [Xe] 4f526S2 69 Tm [Xe] 4f356S2 70 Yb [Xe] 4f546S2 71 Lu [Xe] 4f45d16S2 72 Hf [Xe] 4f45d26S2 73 Ta [Xe] 4f45d36S2 74 W [Xe] 4f45d46S2 75 Re [Xe] 4f45d56S2 76 Os [Xe] 4f45d66S2 77 Ir [Xe] 4f45d76S2 78 Pt [Xe] 4f45d86S1 79 Au [Xe] 4f45d106S1 80 Hg [Xe] 4f45d106S2 81 Ti [Xe] 4f45d106S26P1 82 Pb [Xe] 4f45d106S26P2 83 Bl [Xe] 4f45d106S26P3 84 Po [Xe] 4f45d106S26P4 85 At [Xe] 4f45d106S26P5 86 Rn [Xe] 4f45d106S26P6 or [Rn] 87 Fr [Rn]7S1 88 Ra [Rn]7S2 89 Ac [Rn] 6f57S2 90 Th [Rn] 6f57S2 91 Pa [Rn] 5f46d17S2 92 U [Rn] 5f46d17S2 93 Np [Rn] 5f46d17S2 94 Pu [Rn] 5f57S2 95 Am [Rn] 5f57S2 96 Cm [Rn] 5f46d17S2 97 Bk [Rn] 5f597S2 98 Cf [Rn] 5f5107S2 99 Es [Rn] 5f5117S2 100 Fm [Rn] 5f5127S2 101 Md [Rn] 5f5137S2 102 No [Rn] 5f5147S2 103 Lr [Rn] 5f4146d17S2 27 CHM 101 INTRODUCTION INORGANIC CHEMISTRY 104 Unq [Rn] 5f4146d27S2 105 Unp [Rn] 5f4146d37S2 106 Unh [Rn] 5f4146d47S1 107 Uns [Rn] 5f4146d57S2 108 Uno [Rn] 5f4146d67S2 109 Une [Rn] 5f4146d77S2 3.2 Electronic Configuration of all the Elements in the Periodic Table Period 1 This contains only hydrogen and helium and it is the smallest of all the periods of the table. Hydrogen (Z=1) and helium (Z = 2) are the two elements belonging to the period. The electronic configuration of hydrogen and helium are 1S1 and 1S1 respectively. Thus the 1S orbital which is the only orbital coresponding to 1 is completely filled. The 2S2 configuration of helium is usually represented by [He], so any time you see [He] electron configuration it represents 1S2. Period 2 This period contains elements from lithium (Z=3) to neon (Z=10). In lithium and beryllium, the filling of 2S orbital takes place, then in the next six elements from boron to neon, the 2p orbitals are filled. Neon thus has the electronic configuration of [He] 2S22P6 which as was done in the case of He, is represented by [Ne]. (An electronic configuration of [Ne] means 1S22S22P6. At this stage, the shell having n=2 is complete. Period 3 This is similar to period 2, this period also consists of 8 elements from sodium (Z=11) to argon (Z=18), these elements 3S and 3p orbitals are successively filled in the sequence just as was done in period 2, thus argon has the electronic configuration [Ne] 2S22P6 represented as [Ar] although the third principal shell (n = 3) can accommodate 10 more electron in 3d orbitals, filling of 4s orbital takes place first because of its lower energy but it does not immediately expand from 8 to 18 until scandium is reached Period 4 This period 18 elements from potassium (Z=19) to krypton (Z=36). In K and Ca, the first two elements of this period, the successive electrons go into the 4s orbitals giving them the configuration [Ar] 4S1 and [Ar] 28 CHM 101 MODULE 1 4S2 respectively. Then in the following 10 elements (Sc, Ti, V, Cr Mn, Fe, Co, Ni, Cu and Zn) filling of hither to unoccupied 3d orbitals takes place. Thus the electronic configuration of zinc becomes [Ar] 3d104S2 Occasionally an electron from 4s orbitals is shifted out of turn to the 3d orbitals due to higher stabilityof half-filled and completely filled orbitals, for example, Cr (Z=24) and Cu (Z=29) have the configuration [Ar] 3d54S1 and [Ar] 3d104S1 instead of the· expected [Ar] 3d104S2 and [Ar] 3d94S2 respectively. After the 3d level is filled, in the next six elements of this period, that is Ga, Ge, As, Se, Br and Kr, the 4p orbitals are gradually filled and Kr has the electronic [Ar] 3d104S24P6 represented as [Kr]. Period 5 The next 18 elements from rubidium (Z=37) to Xenon (Z=54) belong to this period. In building up of the atoms of these elements, 5S 4d and 5p orbital are successively filled just as the 4s 3d, Cd, 4p are filled in the elements of period 4. In Rb (Z=37) and Sr (Z=38), the 5S orbital is filled. After that in elements fromy (Z=39) to Cd (Z=48) filling of 4d orbitals takes place. You can see from table 8 that once again there are minor irregularities in the distribution of electron between 4d and 5S orbitals. For example, Mo (Z=42) and Ag (Z=47) have respectively [Kr] 4d55S1 and [Kr] 4d105S1 configurations similar to those of Cr and Cu respectively. Anomalous electronic configuration of Nb, Ru, Rh and Pd cannot be explained in simple terms. You have to therefore, remember them as exceptions. Now in the next six elements, that is I, Sn, Sb, Te, I and Xe filling of S P orbitals take place and thus Xe (Z=54 attains [Kr] 4d105S25P6 configuration Period 6 This period contains 32 elements from caesium (Z=55) to radon (Z=86) in which the 6S, 4f, 5d and 6p orbitals are filled. The first two elements of this period have configurations analogous to those of corresponding member of the lower periods, thus caesium (Z=55) and barium (Z=56) have [Xe] 6S1 and [Xe] 6S2 configuration respectively. According to aufbau principle in the next element La (Z = 57), the additional electron should enter 4f orbital. Instead it goes to the 5d orbitals and La has the configuration [Xe] 5d16S2, but why? The extra electron in the building up of La atom goes to 5d orbital instead of 4f orbital because in La atom, the 5d and 4f orbitals have almost the same energy and hence, the electron is free to enter any of these two orbitals. In the next 14 elements from cerium (Z=58) to lutecium (Z=7l), the 4f orbital is successively filled pertaining to [Xe] 4f15d16S2 and [Xe] 4f144d16S2 configuration, respectively, but you should remember, it is 29 CHM 101 INTRODUCTION INORGANIC CHEMISTRY only Ce (Z=58) Gd (Z=64) and Lu (Z=71) that 5d orbitals have one electron while in all the remaining Lanthanides the 5d orbitals remain vacant after Lutecium, successive electrons occupy 5d orbitals and the electronic configuration builds up from [Xe] 4f144d26S2 for hafnium to [Xe] 4f144d106S2 for mercury the homologue of zinc and cadmium. Again a minor departure from a steady increase in the number of d electrons occurs. For example, gold has [Xe] 4f144d96S2, and as you can see, has to do with the greater stability of half-filled/fully filled orbitals. Finally, the period is completed with successive occupation of the 6p orbitals from thallium, [Xe] 4f145d106S26p1 to radon, [Xe] 4f145d106S26p6. Period 7 This period is still in complete and contains 23 elements from francium (Z=87) to unnitennium (Z =109). In these elements electrons are Filled in 7s, 5f and 6d orbitals. Francium ([Ru] 7S1), radium ([Ru] 7S2) and antinium ([Ru] 6d17S2) have electronic configurations analogous to those of caecium, barium and lanthanum respectively. Thorium has the configuration [Ru] 6d27S2. Therefore, in the 13 elements from protactinium (Z=91) to lawrencium (Z=103) filling of Sf orbitals takes place successively. However, out of these only Pa (Z=91), U (Z=92), Np (Z=93), Cm (Z=96) and Lr (Z=103) have an electron in 6d orbitals. In the rest of the elements, the 6d orbitals remain vacant, thus the electronic configuration of Lr (Z=103) is [Ru] 4f146d27S2 The next six known elements of this period are members of 6d transition series which have the configurations [Ru] 4F145d26p27S2 to [Ru] 4f146d77S2 Having examined the electronic configuration of elements in the periodic table, you can see from table 8 that the elements occupying the same group of the periodic table have the same valence-shell electronic configuration. In order words, the elements having the same valence shell electronic configuration recur periodically, that is after intervals of 2, 8, 8, 18, 18 and 32 in their atomic number. Therefore, periodicity in the properties of elements can easily be understood. In-text Question Question State why Cr (Z=24) and Cu (Z=29) have the configuration [Ar] 3d54S1 and [Ar] 3d10451 instead of the· expected [Ar] 3d104S2 and [Ar] 3d94S2 respectively. 30 CHM 101 MODULE 1 3.3 Electronic Configurationof Ions So far, we have studied the electronic configuration of neutral atoms of elements. I am sure you will be interested in knowing the electronic configuration of ions that are obtained by removal of electrons from the elements. When the gaseous iron atom having [Ar] 3d64S2 ground state electronic configuration loose an electron, the Fe+ ion is formed. These ions have its minimum energy in the configuration [Ar] 3d7, although the iso-electronic manganese atom has the configuration [Ar] 3d54S2 in the ground state. Similarly, the ground state of the Fe2+ and Fe3+ ions are [Ar] 3d6 and [Ar] 3d5 respectively rather than [Ar] 3d54S1 and [Ar] 3d34S2 which are ground states of iso-electronic atoms of chromium and Vanadium respectively. Evidently the differences in nuclear charge between Fe+ and Mn, Fe2+ and Cr and Fe3+ and V are important in determining the orbital to be occupied by the electrons. However, along the series of ions carrying the same charge, the electronic configuration often changes much more regularly than the electronic configuration of the corresponding atoms. Thus for dipositive ions Sc3+ to Zn2+, the ground state electronic configuration changes regularly from [Ar] 3d1 to [Ar] 3d10 for tripositive ions, there is a similar regular change from [Ar] for Sc3+ to [Ar] 3d9 for Zn3+. For tripositve ions of lanthium elements, thereis a regular change from [Xe] 4f1 for Ce3+ to [Xe] 4f14 for Lu3+. Since the chemistry of elements is essentially that of their ions, the regularities in configuration of ions are much more important than the irregularities in the electronic configuration of the neutral atoms. 3.4 Electronic Configuration and Division of Elements into Blocks Elements of the periodic table have been divided into four blocks s, p, d and f depending upon the nature of the atomic orbitals into which the differentiating or the last electron enters. The S-Block Elements-In these elements the differentiating electron enters the 'ns' orbital. Alkali and alkaline earth metals of groups (IA) and 2 (llA) belong to this block. As you know the valence shell electronic configuration of these groups are ns1 and ns2 respectively. We also know that each period of the periodic table begins with alkali metals. All the elements in thi s block are metals. The p-Block Elements- In the elements belonging to this block, the p- orbitals are successively filled. Thus the elements of the group 13 (1llA), 14(IVA), 15(VA), 16(VIA), 17(Vl1A) and 18(zero) are members of this block, since in the atoms of these elements, the differentiating electron enters the np orbitals. The ns orbital in the atoms of these elements are 31 CHM 101 INTRODUCTION INORGANIC CHEMISTRY already completely filled so they have the valence shell electronic configuration ns2np1-6 Note that the elements of s-and p-blocks are also known as normal representative or main group elements. The d-Block Elements- The elements in which the differentiating electron enters the (n-1)d orbitals are called d-block elements. These elements are placed in the middle of the periodic table between the s- and p- block elements. The electronic configuration of the atoms of the elements of this block can be represented by (n-1)d1-10 ns0-12 These elements which are also called transition elements are divided into four series corresponding to the filling of 3d- 4d- 5d- or 6d- orbitals while the 3d, 4d, and 5d series consist of 10 elements each, the 6d series is incomplete and has only seven elements viz: Ac (Z=89) and from Unq 9 (Z=104) to Une (Z=109). The element from Sc(Z=21) to Zn(Z=30),Y(Z=39) to Cd(Z=48), La(Z=57) and from Hf(Z=72) to Hg (Z =80) are the members of 3d, 4d, and 5d series respectively. Note: That D-Block elements are also known as transition elements. The f-Block Elements- The elements in which the extra electron enters (n-2) f orbitals are called the f-block elements. The atoms of these elements have the general configuration (n-2) f1-14(n-l) d0-1ns2. These elements belong to two series depending upon the filling of4f and 5f orbitals. Elements from Ce (Z = 58) to Lu (Z = 71) are the members of the 4f series, while those from the (Z=90) to Lr (Z=103) belong to the 5f series. Elements of 4f series which follow lanthanium in the periodic table are known as LANTHANIDES whereas those of 5f series following actinium are called ACTINIDES. All these elements are collectively referred to as INNER-TRANSITION elements because of filling of electrons in an inner (n-2) f sub shell. Note that f-block elements are also known as inner transition elements. In-text Question Question Explain Hunds rule of maximum multiplicity SELF-ASSESSMENT EXERCISES i. What principles or rules are violated in the following electronic configuration? Write the names of the principle or rule in the space provided along side each configuration. a. 1S22S3 b. 1S22S22Px2 2py1 32 CHM 101 MODULE 1 c. 1S22Px2 ii. Write the electronic configuration of the atoms whose atomic numbersare: a. 21 b. 24 c. 29 4.0 CONCLUSION In conclusion, we have seen that the order in which electrons occupy atomic orbitals is governed by certain rules and principles. These rules determine how many and which electrons occupy valence shells. It is the valence shells that determine the kind of reaction an atom will be involved in. It is therefore very important that the order of filling of orbitals is properly understood. Having examined the electronic configuration of elements in the periodic table, you can see from table 8 that the elements occupying the same group of the periodic table have the same valence- shell electronic configuration. In order words, the elements having the same valence shell electronic configuration recur periodically, that is after intervals of 2, 8, 8, 18, 18 and 32 in their atomic number. Therefore, periodicity in the properties of elements can easily be understood. 5.0 SUMMARY In summary, we have studied the following in this unit: 1. That the filling in of electrons into their orbitals is governed by a) The aufbau principle which assumes that there exist a set of empty hydrogen like orbitals into which electrons can be added b) The n+l rule, which states that in building up electronic configuration of the elements the sub shell with the lowest value of n=1 fills first. c) The Pauli Exclusion principles which states that no two electrons in the same atom can have the same value of all four quantum numbers. d) The Hund's rule which states that as far as possible in a given atom in the ground state, electrons in the same sub shell will occupy different orbitals and will have parallel spins. 2. That the electronic configuration of ions changed regularly. 3. That the elements in the periodic table are divided into four blocks viz: S-P-D- and F Blocks. 33 CHM 101 INTRODUCTION INORGANIC CHEMISTRY 6.0 TUTOR MARK ASSIGNMENTS 1. Explain Pauli's exclusion principle 2. Distinguish LANTHANIDES and ACTINIDES. 7.0 REFERENCES/FURTHER READINGS J. G Wilson, A. B. Newell (1971) General and Inorganic Chemistry (2nd edition) Published by Cambridge University Pres F. A Cotton, G. Wilkinson and P. L. Gayus (1995) Basic Inorganic Chemistry. (3nd edition) John Wiley and Sons Published Gary L. Miessler, Paul J. Fischer and Donald A. Tour (2014) Inorganic Chemistry. 5th edition. Pearson Publishers. ISBN 10:0-321- 81105-4 J. D Lee (2016) Concise Inorganic Chemistry. (5th edition)Wiley Publishers Rules governing the filling of electrons in orbital https://www.youtube.com/watch?v=ZJRcwMtnlAY Electronic Configurationof Ions https://www.youtube.com/watch?v=oCajIGPK-WM Electronic Configuration and Division of Elements into Blocks https://www.youtube.com/watch?v=hJ0WwAGSjlc 34 CHM 101 MODULE 1 UNIT 4 ATOMIC RADII 1.0 Introduction 2.0 Intended Learning Outcomes (ILOs) 3.0 Main content 3.1 Atomic Radii 3.2 Covalent Radius 3.3 Vander Waal's Radius 3.4 Metallic Radius 3.5 Ionic Radius 3.6 Factors affecting atomic radii 3.7 Periodicity in Atomic radii 4.0 Conclusion 5.0 Summary 6.0 Tutor Mark Assignments 7.0 References/ Further Readings 1.0 INTRODUCTION In units 1-3 we studied the development of the periodic law and the periodic table. We learned about the properties of elements being periodic function of their atomic numbers. We learned how electrons are arranged in their orbitals. Arrangements that give rise to similarities and differences in the properties of elements whose valence electrons appear in the same group and those whose valence electrons are in different groups respectively. These differences in the properties arise due to differences in atomic properties. Such as size of the atoms as measured in terms of radii. In this unit, you will be studying about different types of atomic radii, factors affecting atomic radii and periodicity in atomic radii. 2.0 OBJECTIVES By the end you this unit, you should be able to:  Define accurately"Atomic radii"  Distinguished between various forms of atomic radii  List and explain with 80% accuracy the two factors affecting atomic radii  Explain using examples periodicity in atomic radii. 35 CHM 101 INTRODUCTION INORGANIC CHEMISTRY 3.0 MAIN CONTENT 3.1 MEASUREMENT OF ATOMIC RADII Atomic radii are the measure of the size of the atom. Atomic radii are important because other atomic properties like ionization energy, electron affinity and electronegativity are related to them. The wave mechanical picture of an atom depicts an atom as composed of a compact nucleus surrounded by an electron cloud. This electron cloud does not have a definite boundary surface like that of a ball. There is a definite but very small probability of finding an electron at an infinite distance from the nucleus of the atom. However, this does not mean that the atom. is indefinitely large, therefore we have to find away to define the size of an atom. Accordingly, the radius of an atom can be defined as the distance from the centre of the nucleus to the point where the electron density is virtually zero. Now that we have defined the size of an atom, we have to tackle the problem of measuring that size. We are immediately confronted with the problem of defining and accurately measuring the size we mean. Thus, if we are measuring the size of an atom when it is occupying a lattice site of the crystal, the value will be different from one when it is colliding with another atom in the gaseous state. Further more, the size of a neutral atom will be different from the one when it is present as a cation or anion. Consequently, we can not have one set of atomic radii applicable under all conditions. It therefore becomes necessary to specify the bonding conditions under which the size is being measured. Pertaining to the four major types of bonding, the atomic radii to be measured are: 1. Covalent radius 2. Crystal or Metallic radius 3. Vander Waals radius 4. Ionic radius 3.2 Covalent Radius It is possible to determine the distance between the nuclei of bonded atoms by employing Spectroscopic or X-ray and electron diffraction techniques. The covalent or atomic radii are additive that is the bond distance in C-Cl is almost the same as that obtained by adding together the atomic radius of carbon and chlorine Covalent radius can be defined as one half of the distance between the nuclei of two like atoms bonded together by a single covalent bond. 36 CHM 101 MODULE 1 For simplicity simplicity homonuclear diatomic molecules will be examined first, Homonuclear means that there is only one type of nucleus, that is one element present and diatomic means that the molecules are composed of two atoms. If in a homonuclear diatomic molecule of A2 type (eg F2, Cl2, Br2, 12) rA- A is bond length or inter nucleus distance and rA is the covalent radius of the atom A, then rA= 1/2 rA-A. The internuclear distance rC-C between two carbon atoms in diamond is 154pm so the covalent radius of carbon, is equal to77pm. Similarly, the rCl-Cl for solid l2 is 198 pm, rCl is therefore 99pm. In the heteronuclear, diatomic molecule of AB type, if the bonding is purely covalent, then the bond length rA-B is equal to the sum of covalent radii of A and B that is rA-B=rA+rB. Thus covalent radii are additive. It is possible to calculate the radius of one of the atoms in a heteronuclear diatomic molecule of AB type. If we know the internuclear distance rA-B and radius of the other atom. For example, the Si-C bond length in carborundum is 193 pm and covalent radius of C is 77, so you can calculate the covalent radius of Si as follows: rSi - C= rSi+ rC or rSi=rSi-C - rC or rSi=193-77=116pm As stated earlier, the above relatio holds well only ifthe bond between the atoms A and B is purely covalent. If there is a difference inthe electronegativities of the bonded atoms, it causes shortening of the bonds. Schoemaker and Stevenson have proposed the following relationship between the shortening of the bond and the electronegativity difference of the atoms; rA-B = rA + rB – 0.07 (XA –XB)2 Here XA and XB are the electronegativities of A and B respectively. Multiplicity of the bond also causes a shortening of the bond. Usually a double bond is about 0.86 times and a triple bond about 0.78 times the single bond length for the second period elements. Covalent radii of the elements are listed in table 9. In-text Question Question Explain the wave picture of the atom. 37  CHM 101 INTRODUCTION INORGANIC CHEMISTRY Table 9. Covalent and Vander Waals radii of elements 1 2 3 4 5 6 7 8 9 10 11 12 1 IA IIA IIIB IVB VB VI VIIB VIIIB IB IIB I H 37 120 Li Be 123 89 B Na Mg 156 136 A K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn G 103 174 144 132 122 118 117 117 116 115 117 125 1 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd I 216 191 162 145 134 130 127 125 125 128 134 144 1 Cs Ba La* Hf Ta W Re Os Ir Pt Au Hg T 235 198 169 144 134 130 128 126 127 130 134 147 1 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 168 165 164 - 166 185 161 159 159 158 157 156 170 156 38 CHM 101 MODULE 1 3.3 Van Der Waal's Radius In the solid state, non-metallic elements usually exist as aggregates of molecules. The bonding within a non metal molecule is largely covalent. However, individual molecules are held together by weak forces known as Vander Waals forces. Half of the distance between the nuclei of two atoms belonging to two adjacent molecules in a crystal lattice is called Vander Waal's radius. Table 9. Lists the values of Van der Waals radii of some elements. Figure 3. Illustrate the difference between the covalent and vander Waals radii of chlorine Fig 3. Covalent and vander Waals radii of solids chlorine It is evident from the figure that half of the distance between the nuclei X and X1 of the two non-bonded neighbouring chlorine atoms of adjacent molecule A and B is the Vander Waal's radii of chlorine atom. On the other hand, half of the distance between the two nuclei X and Y in the same molecule is the covalent radius of chlorine atom. Thus Van der Waal’s radii represent the distance of the closest approach of an atom to another atom it is in contact with, but not covalently bond to it. Values of Vander Waals radii are larger than those of covalent radii because vander Waals forces are much weaker than the forces operating between atoms in a covalently bonded molecule. 3.4 Metallic or Crystal Radius Metallic or crystal radius is used to describe the size of metal atoms which are usually assumed to be loosely packed spheres in the metallic crystal. The metal atoms are supposed to touch one another in the crystal. Metallic radius is defined as o n e -half of t h e distance between the nuclei of two adjacent metal atoms in the close packed crystal lattice, for example the internuclear distance between two adjacent Na atom in a crystal of sodium metal is 382 pm so metallic radius of Na metal is 382 that is 191pm. 39 CHM 101 INTRODUCTION INORGANIC CHEMISTRY The metallic radius depends to some extent on the crystal structure of the metal. Most metals adopt a hexagonal close packed (hcp) or cubic close packed (ccp) or body centered cubic (bcc) lattice (see figure 4) Fig. 4. Types of metallattices (a) hexagonal; (b) cubicclosepacked (c) body-centredcubic 40  CHM 101 MODULE 1 Table 10. Metalic and ionic radii of elements 1 1 1 1 1 1 1 11 1 2 3 4 5 6 7 8 9 0 1 2 3 4 5 6 78 I V I V II I V V II I II I V V II I II I V V I II I A A B IVB B I B VIIIB B B A A A A A A H 2 0 8( - H 1) e B Li e 1 1 5 1 5 2 6 3 0( 1( + + N 1) 2) B C N O F e N M a g 1 1 9 6 0 0 9 6 5( 5( + + A C A 1) 2) l Si P S l r T C M F C N C S B i V r n e o i u G G A e r 1 1 1 1 1 1 1 1 a e s 1 1 4 3 3 3 2 2 2 2 1 1 1 4 1 C 7 5 0 5 6 5 5 8 Z 4 4 2 0 4 K a 7 8 8 8 7 7 7 6 n 1 1 9 4 3 2 1 6 4 4 0 6 2 2 9( 1 1 1 4 7 9 3 9 ( ( ( ( ( ( ( + 3 1 1 7( ( ( 5 7 + + + + + + + 1) 7 3( 3( + + + 1 9 Sc ) ) ) ) 2 2 2 9 7 + + 5) 6 7 3 9( 164 6 6 6 6 ) ) ) 6( 4( 1) 1) 1 ) ) 3( + 81( 8 6 9 6 6 6 6 + + 6 6 9 1 1 K +) 2) +3) ( ( ( ( 4 3 2 2) 2) 2( 2( 8( 9 9 r 41  CHM 101 INTRODUCTION INORGANIC CHEMISTRY + + ( + + ( ( + + - 8 5 ) ) + ) 3 + + 3) 3) 2) (- (- 3 ) 3 3 2 1 ) ) ) ) ) M R T o u e 1 1 1 3 3 S S 6 9 4 In n b 0 I 6 6 1 1 1 5 5 Z N 8 9 R P 6 6 5 6 0 R S r b ( ( h d A 6 2 9 ( ( b r 1 1 + + 1 1 g C 1 1 6 + + 2 2 6 4 4 3 3 3 1 d 3 1 2( 6 7 4 1 0 6 ) ) 4 7 4 1 2( 2( + ) ) 8 5 8 7 6 6 8 9 4 5 + + 5) 2 2 1 1 0 0 2 T 7 6 6 1 4 1) 2) 2 1 1 4 1 Y ( ( ( c ( ( ( 2 9 8 7 4 1 6 8( 3( 178 + + + 1 + + + 6( 7( 1( 1( 5( (- (- + + 93( 4 5 6 3 4 2 2 + + + + - 2 1 X 1) 2) +3) ) ) ) 6 ) ) ) 1) 2) 3) 4) 3) ) ) e W 1 4 P 1 Ti b Bi 6 1 1 1 H T 8 O I P 7 7 7 C B f a ( s r t A H 1 5 0 s a 1 1 + 1 1 1 u g 1 1 1 2 2 6 4 4 3 3 3 1 1 4 2 2 6 2 0 9 ) 5 4 9 4 5 0( 0( 0( 7 2 La 8 7 6 6 6 9 6 7 + + + 1 1 * 1 3 4 R 9 6 6 1 1 1) 3) 3) P 6 3 188 ( ( ( e ( ( ( 3 1 9 8 7 o 9( 5( 115 + + + 1 + + + 7( 0( 5( 4( 4( 1 + + (+3 4 5 6 3 4 4 2 + + + + + 7 A R 1) 2) ) ) ) ) 7 ) ) ) 1) 2) 3) 4) 5) 6 t n 42 CHM 101 MODULE 1 In both these structure, a given metal atom has twelve nearest neighbours. However, a significant number of metals adopt a body centred cubic lattice (BCC) in which the number of nearest neighbours is eight. The number of nearest neighbours of a metal atom in a lattice is known as the coordination number of the metal. Experimental studies on a number of metals having more than one crystal lattice have shown that the radius of a metal in an eight coordinate lattice is about 0.97 of the radius of the same metal in a twelve coordinate environment. These bonds, eight or twelve per metal atom. On the other hand, the metallic crystal lattices are stronger than the Vander Waals forces. Table 10. Gives a set of twelve coordinate radii for metal atoms. Compare these with the covalent radii or Vander Waals radii in table 9. The metallic radii are generally larger than the corresponding covalent radii. Although both involve a sharing of electrons this is because the average bond order of an individual metal-metal bond is considerably less than one and therefore the individual bond is weaker and longer than the covalent. This does not mean that the overall bonding is weak as there is a large number of hese bonds, eight or twelve per metal atom. On the other hand, the metallic crystal lattices are stronger than the Van der waals forces. 3.5 Ionic Radius Ionic radius is defined as the distance between the nucleus of anion and the point up to which the nucleus has influence on the electron cloud. In order words, it may also be defined as the distance of the closest approach from the centre of ion by another ion. Ionic radius is usually evaluated from the distance determined experimentally between the centres of nearest neighbors. Thus if we wish to estimate the ionic radius of Na+ we may measure the internuclear distance between Na+ and Cl- ions in the NaCl crystal lattice. This distance is the sum of radii of Na+ and Cl- ions. From the electron density maps obtained by X-ray analysis, it has become possible, in some cases, to apportion the internuclear distance into ther adius of cation and anion. A small member of ionic crystals has thus been studied and the ionic radii of some of the elements have been determined. These radii have become the basis for assigning the ionic radii of most of the other elements. Values for ionic radii cannot be measured absolutely, but are estimated. They are not completely accurate or reliable. Though it is possible to measure the interatomic distance between two different ions very accurately by X-ray crystallography. Ionic radii are of two types, cation radii and anion radii. All common cations are smaller than all common anion except for rubidium and caesium cations (largest single atom cations). This is not too surprising since not only is there a loss of electron(s) from apartially filled outer 43 CHM 101 INTRODUCTION INORGANIC CHEMISTRY shell on cation foarmaton, but there is also an increase in the over all positive charge on the ion. Conversely, in anion formation the addition of an electron to an atom increases the size due to increase in inter-electronic repulsion in the valence-shell and decrease in effective nuclear charge. In general, there is a decrease in size of anions to covalent radii of corresponding atoms to cations thus in the series of isoelectronic species (eg. N3-, O2-, Ne, Na+, Mg2+ and Al3+). The greater the effective nuclear charge, the smaller is the radius of the species. In table 10; radii of some of the common ions have been listed. 3.6 Factors Affecting the Atomic Radii So far, we have defined and explained types of atomic radii. We shall now turn our attention to two of the factors that affect them. 1. (Principal QuantumNumber (n): The principal quantum n has integral values of 1, 2, 3, 4, etc. As the principal quantum number increases, the outer electrons get farther away from the nucleus and hence the atomic radius generally increases, been with largest value of n has the most energy, therefore it requires the least input of energy to ionize it. 2. Effective Nuclear Charge (Z*): The magnitude of the effective nuclear charge determines the·magnitude of the force of attraction exerted by the nucleus on the outer most electrons. The greater the magnitude of effective nuclear charge, the greater is the force exerted by the nucleus on the outermost electron. Hence theelectron cloud of the outer most shell is pulled inward nearer to the nucleus and consequently its distance from the nucleus. That is, atomic radius decreases. Effective nuclear charge Z* is the amount of positive charge felt by the outer electrons in an atom. It is always less than the actual charger Z of the nucleus of the atom. This is because electrons in inner shells partially shield the electrons in the outer shell from nuclear attraction. The effective nuclear charge felt by the outer electron depends upon the actual nuclear charge and the number and type of inner screening electrons. It can be calculated by subtracting he screening or shielding constant, S from the atomic number Z Thus Z* =Z-S. You can estimate the value of screening constant, S, with the help of Slater’s rules in the following manner: 44 CHM 101 MODULE 1 1. Write out the electronic configuration of the element in the following order and groupings; (ls) (2s, 2p) (3s, 3p) (3d) (4s, 4p) (4d) (4f) (5s, 5p) (5d) (5f) ((6s, 6p) etc. 2. Electrons in any group higher in this sequence than the electron under consideration contribute nothing to s. For example, in Ti atom (electronic configuration IS2 2S2 2p6 3S2 3p6 3d2 4S2). The two electrons in 4S orbital will contribute nothing towards the screening for an electron in3d orbital. 3. Then for an electron in an ns or np orbitals (a) All other electrons in the (ns, np) group contribute S=0.35 each except for the electron in is which contribute S=0.30 (b) All electron sin (n-1) shells contribute S=0.85 each (c) All elctrons in (n-2) or lower shells contribute S=1.00 each 4. For an electron in an nd or nf orbrtal (a) All electrons in the same group that is nd or nf contribute S= 0.35each. (b) Those in the groups lying lower in the sequence than the nd or nf group contribute S=1.00 each. In order to demonstrate the application of Slater's rules, we shall now calculate the Z*for an electron in N, K and Zn atoms. 1. Electronic configuration of N= IS2 2S2 2p3 Grouping (IS2) (2S2 2p3) Value of screening constant for an electron in 2p orbital will be S= (4x0.35) + (2x0.85) =3.10. Hence Z*=Z-S=7-3.10=3.90 2. Electronic configuration of K=1S22S22p63S23p64s1 Grouping of orbitals will be (IS2) (2S22p6) (3S2 3p6) (4s1) Value of screening constant for an electron in 4s orbitals will be S=90.85 x8) + (Ix10) =16.80. Hence effective nuclear charge Z*=Z- S=19-16.80= 2.20 3. Electronic configuration of Zn= 1S22S22p63S23p63d10 4S2 Grouping of the orbitals gives (1S2) (2S22p6) (3S23p6) (3d10) (4S2) Value of screening constant for S for an electron in 4S orbital will be S= (0.35x1) + (0.85 x18) + (1x10) = 25.65 hence the effective nuclear charge felt by 4S electron will be Z* = Z- S = 30-25.65=4.35 If we consider a 3d electron m Zn the grouping is as above, but the effective nuclear charge felt by the 3d electron will be Z* =Z- S=30- [(9 45 CHM 101 INTRODUCTION INORGANIC CHEMISTRY x0.35) = (18 x1)] =8.85. Thus you can see an electron in 3d orbitals in Zn is more strongly old by the nucleus than that in 4S orbital Table 11 contains a list of values of nuclear charge for electron in valence shell in the first thirty elements calculated by Slater's rules. You can see from the table that there is a steady increase in Slater's Z* across rows of the periodic table. Effective nuclear charge felt by electrons also depends on the oxidation state of an atom in a compound. The higher the oxidation state of the atom, the higher will be the effective nuclear charge felt by the electrons and therefore, smaller will be the atomic radius. Thus the ionic radius of Fe3+ ion will be smaller than that of the Fe2+ ion. Similarly, covalent radius of bromine in BrCl, will be then that in BrCl. 46 CHM 101 MODULE 1 Table 11: Effectiveness of nuclear charge for first 30 elements Period Element Z S Z* 1 H 1 0 1.0 2 He 2 0.3 1.70 Li 3 1.70 1.30 Be 4 2.05 1.95 B 5 2.40 2.60 C 6 2.75 3.25 N 7 3.10 3.90 O 8 3.45 4.55 F 9 3.80 5.20 Ne 10 4.15 5.85 3 Na 11 8.80 2.20 Mg 12 9.15 2.85 Al 13 9.50 3.50 Si 14 9.85 4.15 P 15 10.20 4.80 S 16 10.55 5.45 Cl 17 10.90 6.10 Ar 18 11.25 6.75 4 K 19 18.80 2.20 Ca 20 17.13 2.85 Sc 21 18.0 3.0 Ti 22 18.85 3.15 V 23 19.70 3.30 Cr 24 20.55 3.45 Mn 25 21.40 3.60 Fe 26 22.25 3.75 Co 27 23.10 3.90 Ni 28 23.95 4.05 Cu 29 24.80 4.20 Zn 30 25.65 4.35 3.7 Periodicity in Atomic Radii Now that we know the various types of atomic radii and the factors that affect them, we will consider the periodicity in them. Before doing that however, we would like to emphasize that trends observed in one type of radii (example covalent radii) are generally found in the other type of radii also (example ionic and metallic radii). Two general periodic trends are found for all types of atomic radii. These are the atomic radii decreases along a period and generally increase down a group in the long form of the periodic table (see fig 4). These changes in the atomic radii 47 CHM 101 INTRODUCTION INORGANIC CHEMISTRY can be related to the charges in effective nuclear charge and the principal quantum number in the periodic table. If you examine table 11 you will find out that there is a steady increase (by 0.65 units) in the value of Z* from alkali metals to halogens for the elements of period 2 and 3, but there is no change in the value of n because the electrons fill the same principal shell. As a result of this there is a steady decrease in the covalent radius from 123 and 165pm for Li and Na to 64 and 99pm for F and Cl respectively. In comparison to the above, the decrease in covalent radii across the transition series is much smaller. As you know, electrons are successfully filled in the (n-l) d orbitals across a transition series and hence screen the size determining ns electrons from the nuclear charge more effectively. Therefore, across a transition series, there is only small increase in effective nuclear charge (by 0.15units), therefore only a small increase in effective nuclear charge decrease in atomic radius from one element to another take splace. In 3d series, covalent radius decreases from 144pm for Sc to115pm for Ni. Then in copper and zinc due to completion of 3d sub shell, the electronic charge density in this sub shell becomes very high which increases the inter electronic repulsion. As a result, covalent radii of Cu and Zn increases lightly to 117 and 125pm respectively. Thus across the ten elements of the first transition series, there is an over all decrease in covalent radius by 19pm which is much less than that across seven normal elements of period 2 (59pm) and period 3(57pm). But due to this, the covalent radii of elements from Ga to Kr following Zn becomes much smaller than that expected by simple extrapolation of the values for elements of period 2 and 3 for example, the covalent radii of Al and Ga are equal where as the covalent radii of elements Ge, As, Se, Br are only slightly larger than those of corresponding elements (Si, P, and Cl) of period 3.The rate of decrease in the size across the Lanthanide series is even less than that across the first transition series. In the Lanthanide elements, filling of (n-2)f orbitals take place, while simultaneously the nuclear charge increases. The electrons in the (n-2)f orbital shield then s electrons, (which largely determine the size, from the increase in nuclear charge) almost completely (S=1.00). As a result of this, there is only a small decrease in the atomic radius from one element to another. But there are 14 elements in the series. There is a total of contraction of 13pm across the series from Ca (Z=57) to Lu (Z =71). This is known as lanthanide contraction, because of which the atom s of elements (Hf to Hg) following Lu are usually smaller than they would be if the lanthanide had not been built up before them. Lanthanide contraction almost exactly cancel out the effect of the last shell added in 48 CHM 101 MODULE 1 the sixth period and therefore, the transition elements or 4d and 5d series have almost the same atomic radii. On descending any group of the periodic table, the number of electron in the valence shell remains constant but the number of shells around nucleus increases monotonically, so that the effective nuclear charge felt by valence electrons stays nearly the same. So with increase in principal quantum number (n) of the valence shell, an increase in atomic radii is generally observed down any group of the periodic table. For example, as shown in figure 3, there is an increase in atomic radii of alkali and alkaline earth metals as we proceed downward in the group. However as pointed out earlier, with the inclusion of 3d transition elements in period 4 increase in the radii of elements from Ga to Br is smaller than expected. Similarly, because of inclusion of Lanthanide elements in period 6, atoms of the transition elements of this period (Hf to Hg) are almost of the same size as atoms above than in period 5 (Zr to Cd). After that, only a small increase in size of elements of period 6 (Tc to Al) as compared to the size of elements above them in period 5 (In to I) is observed. Fig 5. Plot of covalent radius against atomic number In-text Question Question Explain the effect of completely filled d-orbital in Cu and Zn. SELF-ASSESSMENT EXERCISES i. Assuming that the atoms are touching each other, what would be the internuclear distance between two fluorine atoms in F2? 49 CHM 101 INTRODUCTION INORGANIC CHEMISTRY ii. Arrange the following isoelectronic species in order of decreasing atomic radius. Na+, Mg2+, Al3+, Si4+, N3-, O2-, F-, Ne 4.0 CONCLUSION Having gone through this unit, we can conclude that knowledge of the size of an atom is indeed very essential. It is through the knowledge of the atomic radii that we can predict accurately their action of the atom. Conversely, in anion formation the addition of an electron to an atom increases the size due to increase in inter-electronic repulsion in the valence-shell and decrease in effective nuclear charge. The bonding within a non metal molecule is largely covalent. However, individual molecules are held together by weak forces known as Vander Waals forces 5.0 SUMMARY In this unit, you have studied the following: i. The definition of atomic radii. ii. The various types of atomic radii viz: covalent, van der Waals metallic and ionic radii iii. Factors affecting radii that is the principal quantum number and the effective nuclear charge Z* iv. Periodicity in atomic radii nuclear chargeZ* v. Periodicity in atomic radii 6.0 TUTOR MARK ASSIGNEMENTS i. a) How doe’s atomic size varies in a group and in a period? Give reasons for the variation. b) Arrange H2, H+ and H- in order of increasing atomic radius ii. What are iso-electronic ions? How does their size vary with the change of atomic number? 7.0 REFERENCES AND FURTHER READING J. G Wilson, A. B. Newell (1971) General and Inorganic Chemistry (2nd edition) Published by Cambridge University Pres F. A Cotton, G. Wilkinson and P. L. Gayus (1995) Basic Inorganic Chemistry. (3nd edition) John Wiley and Sons Published Gary L. Miessler, Paul J. Fischer and Donald A. Tour (2014) Inorganic Chemistry. 5th edition. Pearson Publishers. ISBN 10:0-321- 81105-4 50 CHM 101 MODULE 1 J. D Lee (2016) Concise Inorganic Chemistry. (5th