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70 CHEMISTRY UNIT 3 CLASSIFICATION OF ELEMENTS AND d PERIODICITY IN PROPERTIES...

70 CHEMISTRY UNIT 3 CLASSIFICATION OF ELEMENTS AND d PERIODICITY IN PROPERTIES he The Periodic Table is arguably the most important concept in pu T chemistry, both in principle and in practice. It is the everyday is support for students, it suggests new avenues of research to re ER After studying this Unit, you will be able to professionals, and it provides a succinct organization of the whole of chemistry. It is a remarkable demonstration of the bl fact that the chemical elements are not a random cluster of app reciate how the concept of entities but instead display trends and lie together in families. grouping elements in accordance to An awareness of the Periodic Table is essential to anyone who their properties led to the wishes to disentangle the world and see how it is built up development of Periodic Table. be C from the fundamental building blocks of the chemistry, the understand the Periodic Law; chemical elements. understand the significance of atomic number and electronic Glenn T. Seaborg o N configuration as the basis for periodic classification; name the elements with In this Unit, we will study the historical development of the Z >100 according to IUPAC Periodic Table as it stands today and the Modern Periodic nomenclature; Law. We will also learn how the periodic classification © classify elements into s, p, d, f follows as a logical consequence of the electronic blocks and learn their main configuration of atoms. Finally, we shall examine some of characteristics; the periodic trends in the physical and chemical properties recognise the periodic trends in of the elements. physical and chemical properties of elements; 3.1 WHY DO WE NEED TO CLASSIFY ELEMENTS ? compare the reactivity of elements and correlate it with their We know by now that the elements are the basic units of all occurrence in nature; types of matter. In 1800, only 31 elements were known. By explain the relationship between 1865, the number of identified elements had more than tt ionization enthalpy and metallic doubled to 63. At present 114 elements are known. Of character; them, the recently discovered elements are man-made. use scientific vocabulary Efforts to synthesise new elements are continuing. With appropriately to communicate ideas such a large number of elements it is very difficult to study no related to certain important individually the chemistry of all these elements and their properties of atoms e.g., atomic/ innumerable compounds individually. To ease out this ionic radii, ionization enthalpy, problem, scientists searched for a systematic way to electron gain enthalpy, electronegativity, valence of organise their knowledge by classifying the elements. Not only that it would rationalize known chemical facts about element s. elements, but even predict new ones for undertaking further study. CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 71 3.2 GENESIS OF PERIODIC chemist, John Alexander Newlands in 1865 CLASSIFICATION profounded the Law of Octaves. He arranged the elements in increasing order of their atomic Classification of elements into groups and weights and noted that every eighth element development of Periodic Law and Periodic had properties similar to the first element Table are the consequences of systematising (Table 3.2). The relationship was just like every the knowledge gained by a number of scientists eighth note that resembles the first in octaves through their observations and experiments. of music. Newlands’s Law of Octaves seemed d The German chemist, Johann Dobereiner in to be true only for elements up to calcium. early 1800’s was the first to consider the idea Although his idea was not widely accepted at of trends among properties of elements. By that time, he, for his work, was later awarded he 1829 he noted a similarity among the physical Davy Medal in 1887 by the Royal Society, and chemical properties of several groups of London. three elements (Triads). In each case, he pu T noticed that the middle element of each of the The Periodic Law, as we know it today owes is Triads had an atomic weight about half way its development to the Russian chemist, Dmitri between the atomic weights of the other two Mendeleev (1834-1907) and the German re ER (Table 3.1). Also the properties of the middle chemist, Lothar Meyer (1830-1895). Working bl element were in between those of the other independently, both the chemists in 1869 Table 3.1 Dobereiner’s Triads Element Atomic Element Atomic Element Atomic weight weight weight be C Li 7 Ca 40 Cl 35.5 Na 23 Sr 88 Br 80 o N K 39 Ba 137 I 127 two members. Since Dobereiner’s relationship, proposed that on arranging elements in the referred to as the Law of Triads, seemed to increasing order of their atomic weights, © work only for a few elements, it was dismissed similarities appear in physical and chemical as coincidence. The next reported attempt to properties at regular intervals. Lothar Meyer classify elements was made by a French plotted the physical properties such as atomic geologist, A.E.B. de Chancourtois in 1862. He volume, melting point and boiling point arranged the then known elements in order of against atomic weight and obtained a increasing atomic weights and made a periodically repeated pattern. Unlike cylindrical table of elements to display the Newlands, Lothar Meyer observed a change in periodic recurrence of properties. This also did length of that repeating pattern. By 1868, not attract much attention. The English Lothar Meyer had developed a table of the tt Table 3.2 Newlands’ Octaves Element Li Be B C N O F no At. wt. 7 9 11 12 14 16 19 Element Na Mg Al Si P S Cl At. wt. 23 24 27 29 31 32 35.5 Element K Ca At. wt. 39 40 72 CHEMISTRY elements that closely resembles the Modern weights, thinking that the atomic Periodic Table. However, his work was not measurements might be incorrect, and placed published until after the work of Dmitri the elements with similar properties together. Mendeleev, the scientist who is generally For example, iodine with lower atomic weight credited with the development of the Modern than that of tellurium (Group VI) was placed Periodic Table. in Group VII along with fluorine, chlorine, bromine because of similarities in properties While Dobereiner initiated the study of (Fig. 3.1). At the same time, keeping his d periodic relationship, it was Mendeleev who was responsible for publishing the Periodic primary aim of arranging the elements of Law for the first time. It states as follows : similar properties in the same group, he he proposed that some of the elements were still The properties of the elements are a undiscovered and, therefore, left several gaps periodic function of their atomic in the table. For example, both gallium and weights. germanium were unknown at the time pu T Mendeleev published his Periodic Table. He left is Mendeleev arranged elements in horizontal the gap under aluminium and a gap under rows and vertical columns of a table in order re ER of their increasing atomic weights in such a silicon, and called these elements Eka- bl Aluminium and Eka-Silicon. Mendeleev way that the elements with similar properties occupied the same vertical column or group. predicted not only the existence of gallium and Mendeleev’s system of classifying elements was germanium, but also described some of their more elaborate than that of Lothar Meyer’s. general physical properties. These elements were discovered later. Some of the properties be C He fully recognized the significance of periodicity and used broader range of physical predicted by Mendeleev for these elements and and chemical properties to classify the those found experimentally are listed in o N elements. In particular, Mendeleev relied on Table 3.3. the similarities in the empirical formulas and properties of the compounds formed by the The boldness of Mendeleev’s quantitative elements. He realized that some of the elements predictions and their eventual success made him and his Periodic Table famous. © did not fit in with his scheme of classification if the order of atomic weight was strictly Mendeleev’s Periodic Table published in 1905 followed. He ignored the order of atomic is shown in Fig. 3.1. Table 3.3 Mendeleev’s Predictions for the Elements Eka-aluminium (Gallium) and Eka-silicon (Germanium) Property Eka-aluminium Gallium Eka-silicon Germanium (predicted) (found) (predicted) (found) tt Atomic weight 68 70 72 72.6 Density / (g/cm3) 5.9 5.94 5.5 5.36 no Melting point /K Low 302.93 High 1231 Formula of oxide E2O3 Ga2O 3 EO2 GeO2 Formula of chloride ECl3 GaCl 3 ECl4 GeCl4  CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES PERIODIC SYSTEM OF THE ELEMENTS IN GROUPS AND SERIES h ed pu T is re ER bl be C o N © tt Fig. 3.1 Mendeleev’s Periodic Table published earlier 73 no 74 CHEMISTRY d Dmitri Mendeleev was born in Tobalsk, Siberia in Russia. After his father’s death, the family moved to St. Petersburg. He received his he Master’s degree in Chemistry in 1856 and the doctoral degree in 1865. He taught at the University of St.Petersburg where he was appointed Professor of General Chemistry in 1867. Preliminary work for his great textbook “Principles of Chemistry” led Mendeleev to pu T propose the Periodic Law and to construct his Periodic Table of is elements. At that time, the structure of atom was unknown and re ER Mendeleev’s idea to consider that the properties of the elements bl were in someway related to their atomic masses was a very Dmitri Ivanovich imaginative one. To place certain elements into the correct group from Mendeleev the point of view of their chemical properties, Mendeleev reversed the (1834-1907) order of some pairs of elements and asserted that their atomic masses were incorrect. Mendeleev also had the foresight to leave gaps in the Periodic Table for be C elements unknown at that time and predict their properties from the trends that he observed among the properties of related elements. Mendeleev’s predictions were proved to be astonishingly correct when these elements were discovered later. o N Mendeleev’s Periodic Law spurred several areas of research during the subsequent decades. The discovery of the first two noble gases helium and argon in 1890 suggested the possibility that there must be other similar elements to fill an entire family. This idea led Ramsay to his successful search for krypton and xenon. Work on the radioactive decay © series for uranium and thorium in the early years of twentieth century was also guided by the Periodic Table. Mendeleev was a versatile genius. He worked on many problems connected with Russia’s natural resources. He invented an accurate barometer. In 1890, he resigned from the Professorship. He was appointed as the Director of the Bureau of Weights and Measures. He continued to carry out important research work in many areas until his death in 1907. You will notice fr om the modern Period Table (Fig. 3.2) that Mendeleev’s name has been immortalized by naming the element with atomic number 101, as Mendelevium. This name was pr oposed by American scientist Glenn T. Seaborg, the discoverer of this element, tt “in recognition of the pioneering role of the great Russian Chemist who was the first to use the periodic system of elements to predict the chemical properties of undiscovered elements, a principle which has been the key to the discovery of nearly all the transuranium elements”. no CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 75 3.3 MODERN PERIODIC LAW AND THE Numerous forms of Periodic Table have PRESENT FORM OF THE PERIODIC been devised from time to time. Some forms TABLE emphasise chemical reactions and valence, We must bear in mind that when Mendeleev whereas others stress the electronic developed his Periodic Table, chemists knew configuration of elements. A modern version, nothing about the internal structure of atom. the so-called “long form” of the Periodic Table However, the beginning of the 20 th century of the elements (Fig. 3.2), is the most convenient witnessed profound developments in theories and widely used. The horizontal rows (which d about sub-atomic particles. In 1913, the Mendeleev called series) are called periods and English physicist, Henry Moseley observed the vertical columns, groups. Elements having he regularities in the characteristic X-ray spectra similar outer electronic configurations in their of the elements. A plot of ν (where ν is atoms are arranged in vertical columns, frequency of X-rays emitted) against atomic referred to as groups or families. According to the recommendation of International Union pu T number (Z ) gave a straight line and not the is of Pure and Applied Chemistry (IUPAC), the plot of ν vs atomic mass. He thereby showed groups are numbered from 1 to 18 replacing re ER that the atomic number is a more fundamental the older notation of groups IA … VIIA, VIII, IB bl property of an element than its atomic mass. … VIIB and 0. Mendeleev’s Periodic Law was, therefore, There are altogether seven periods. The accordingly modified. This is known as the period number corresponds to the highest Modern Periodic Law and can be stated as : principal quantum number (n) of the elements be C The physical and chemical properties in the period. The first period contains 2 of the elements are periodic functions elements. The subsequent periods consists of of their atomic numbers. 8, 8, 18, 18 and 32 elements, respectively. The o N seventh period is incomplete and like the sixth The Periodic Law revealed important period would have a theoretical maximum (on analogies among the 94 naturally occurring elements (neptunium and plutonium like the basis of quantum numbers) of 32 elements. actinium and protoactinium are also found in In this form of the Periodic Table, 14 elements © pitch blende – an ore of uranium). It stimulated of both sixth and seventh periods (lanthanoids renewed interest in Inorganic Chemistry and and actinoids, respectively) are placed in has carried into the present with the creation separate panels at the bottom*. of artificially produced short-lived elements. 3.4 NOMENCLATURE OF ELEMENTS WITH You may recall that the atomic number is ATOMIC NUMBERS > 100 equal to the nuclear charge (i.e., number of protons) or the number of electrons in a neutral The naming of the new elements had been atom. It is then easy to visualize the significance traditionally the privilege of the discoverer (or of quantum numbers and electronic discoverers) and the suggested name was tt configurations in periodicity of elements. In ratified by the IUPAC. In recent years this has fact, it is now recognized that the Periodic Law led to some controversy. The new elements with is essentially the consequence of the periodic very high atomic numbers are so unstable that variation in electronic configurations, which only minute quantities, sometimes only a few no indeed determine the physical and chemical atoms of them are obtained. Their synthesis properties of elements and their compounds. and characterisation, therefore, require highly * Glenn T. Seaborg’s work in the middle of the 20t h century starting with the discovery of plutonium in 1940, followed by those of all the transuranium elements from 94 to 102 led to reconfiguration of the periodic table placing the actinoids below the lanthanoids. In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. Element 106 has been named Seabor gium (Sg) in his honour.  76 h ed pu T is re ER bl be C o N © CHEMISTRY tt Fig. 3.2 Long form of the Periodic Table of the Elements with their atomic numbers and ground state outer electronic configurations. The groups are numbered 1-18 in accordance with the 1984 IUPAC recommendations. This notation replaces the old numbering scheme of IA–VIIA, VIII, IB–VIIB and 0 for no the elements. CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 77 sophisticated costly equipment and laboratory. which make up the atomic number and “ium” Such work is carried out with competitive spirit is added at the end. The IUPAC names for only in some laboratories in the world. elements with Z above 100 are shown in Scientists, before collecting the reliable data on Table 3.5. the new element, at times get tempted to claim Table 3.4 Notation for IUPAC Nomenclature for its discovery. For example, both American of Elements and Soviet scientists claimed credit for discovering element 104. The Americans Digit Name Abbreviation d named it Rutherfordium whereas Soviets 0 nil n named it Kurchatovium. To avoid such 1 un u he problems, the IUPAC has made 2 bi b recommendation that until a new element’s 3 tri t discovery is proved, and its name is officially 4 quad q pu T recognized,,,,,,, a systematic nomenclature be 5 pent p is derived directly from the atomic number of the 6 hex h element using the numerical roots for 0 and 7 sept s re ER numbers 1-9. These are shown in Table 3.4. 8 oct o bl The roots are put together in order of digits 9 enn e Table 3.5 Nomenclature of Elements with Atomic Number Above 100 be C Atomic Name according to Symbol IUPAC IUPAC Number IUP AC nomenclature Official Name Symbol 101 Unnilunium Unu Mendelevium Md o N 102 Unnilbium Unb Nobelium No 103 Unniltrium Unt Lawrencium Lr 104 Unnilquadium Unq Rutherfordium Rf © 105 Unnilpentium Unp Dubnium Db 106 Unnilhexium Unh Seaborgium Sg 107 Unnilseptium Uns Bohrium Bh 108 Unniloctium Uno Hassium Hs 109 Unnilennium Une Meitnerium Mt 110 Ununnillium Uun Darmstadtium Ds 111 Unununnium Uuu Rontgenium Rg tt 112 Ununbium Uub Copernicium Cn 113 Ununtrium Uut * – 114 Ununquadium Uuq Flerovium Fl no 115 Ununpentium Uup * – 116 Ununhexium Uuh Livermorium Lv 117 Ununseptium Uus * – 118 Ununoctium Uuo * – * Official IUPAC name yet to be announced 78 CHEMISTRY Thus, the new element first gets a be readily seen that the number of elements in temporary name, with symbol consisting of each period is twice the number of atomic three letters. Later permanent name and orbitals available in the energy level that is symbol are given by a vote of IUPAC being filled. The first period (n = 1) starts with representatives from each country. The the filling of the lowest level (1s) and therefore permanent name might reflect the country (or has two elements — hydrogen (ls1) and helium state of the country) in which the element was (ls2) when the first shell (K) is completed. The discovered, or pay tribute to a notable scientist. second period (n = 2) starts with lithium and d As of now, elements with atomic numbers up the third electron enters the 2s orbital. The next to 118 have been discovered. Official names of element, beryllium has four electrons and has 2 2 elements with atomic numbers 113, 115, 117 the electronic configuration 1s 2s. Starting he and 118 are yet to be announced by IUPAC. from the next element boron, the 2p orbitals are filled with electrons when the L shell is Problem 3.1 completed at neon (2s 22p 6). Thus there are pu T What would be the IUPAC name and 8 elements in the second period. The third is symbol for the element with atomic period (n = 3) begins at sodium, and the added number 120? electron enters a 3s orbital. Successive filling re ER of 3s and 3p orbitals gives rise to the third Solution bl period of 8 elements from sodium to argon. The From Table 3.4, the roots for 1, 2 and 0 fourth period (n = 4) starts at potassium, and are un, bi and nil, respectively. Hence, the the added electrons fill up the 4s orbital. Now symbol and the name respectively are Ubn you may note that before the 4p orbital is filled, and unbinilium. filling up of 3d orbitals becomes energetically be C favourable and we come across the so called 3d transition series of elements. This starts 3.5 ELECTRONIC CONFIGURATIONS OF from scandium (Z = 21) which has the electronic o N ELEMENTS AND THE PERIODIC 1 2 configuration 3d 4s. The 3d orbitals are filled TABLE at zinc (Z=30) with electronic configuration In the preceding unit we have learnt that an 10 2 3d 4s. The fourth period ends at krypton electron in an atom is characterised by a set of with the filling up of the 4p orbitals. Altogether © four quantum numbers, and the principal we have 18 elements in this fourth period. The quantum number (n ) defines the main energy fifth period (n = 5) beginning with rubidium is level known as shell. We have also studied similar to the fourth period and contains the about the filling of electrons into different 4d transition series starting at yttrium subshells, also referred to as orbitals (s, p, d, (Z = 39). This period ends at xenon with the f ) in an atom. The distribution of electrons into filling up of the 5p orbitals. The sixth period orbitals of an atom is called its electronic (n = 6) contains 32 elements and successive configuration. An element’s location in the electrons enter 6s, 4f, 5d and 6p orbitals, in Periodic Table reflects the quantum numbers the order — filling up of the 4f orbitals begins tt of the last orbital filled. In this section we will with cerium (Z = 58) and ends at lutetium observe a direct connection between the (Z = 71) to give the 4f-inner transition series electronic configurations of the elements and which is called the lanthanoid series. The the long form of the Periodic Table. seventh period (n = 7) is similar to the sixth no (a) Electronic Configurations in Periods period with the successive filling up of the 7s, 5f, 6d and 7p orbitals and includes most of The period indicates the value of n for the the man-made radioactive elements. This outermost or valence shell. In other words, period will end at the element with atomic successive period in the Periodic Table is number 118 which would belong to the noble associated with the filling of the next higher gas family. Filling up of the 5f orbitals after principal energy level (n = 1, n = 2, etc.). It can actinium (Z = 89) gives the 5f-inner transition CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 79 series known as the actinoid series. The 4f- theoretical foundation for the periodic and 5f-inner transition series of elements classification. The elements in a vertical column are placed separately in the Periodic Table to of the Periodic Table constitute a group or maintain its structure and to preserve the family and exhibit similar chemical behaviour. principle of classification by keeping elements This similarity arises because these elements with similar properties in a single column. have the same number and same distribution of electrons in their outermost orbitals. We can Problem 3.2 classify the elements into four blocks viz., d How would you justify the presence of 18 s-block, p-block, d-block and f-block elements in the 5th period of the Periodic depending on the type of atomic orbitals that he Table? are being filled with electrons. This is illustrated Solution in Fig. 3.3. We notice two exceptions to this When n = 5, l = 0, 1, 2, 3. The order in categorisation. Strictly, helium belongs to the s-block but its positioning in the p-block along pu T which the energy of the available orbitals with other group 18 elements is justified is 4d, 5s and 5p increases is 5s < 4d < 5p. The total number of orbitals available are because it has a completely filled valence shell re ER 9. The maximum number of electrons that (1s 2) and as a result, exhibits properties bl can be accommodated is 18; and therefore characteristic of other noble gases. The other 18 elements are there in the 5th period. exception is hydrogen. It has only one s-electron and hence can be placed in group 1 (b) Groupwise Electronic Configurations (alkali metals). It can also gain an electron to achieve a noble gas arrangement and hence it be C Elements in the same vertical column or group have similar valence shell electronic can behave similar to a group 17 (halogen configurations, the same number of electrons family) elements. Because it is a special case, o N in the outer orbitals, and similar properties. we shall place hydrogen separately at the top For example, the Group 1 elements (alkali of the Periodic Table as shown in Fig. 3.2 and 1 metals) all have ns valence shell electronic Fig. 3.3. We will briefly discuss the salient configuration as shown below. features of the four types of elements marked in © Atomic number Symbol Electronic configuration 3 Li 1s2 2s1 (or) [He]2s1 11 Na 1s2 2s22p 63s1 (or) [Ne]3s1 19 K 1s2 2s22p 63s2 3p6 4s1 (or) [Ar]4 s1 37 Rb 1s2 2s22p 63s2 3p6 3d104s2 4p6 5s1 (or) [Kr]5s1 55 Cs 1s2 2s22p 63s23p 63d104s2 4p 64d105s2 5p 6 6s 1 (or) [Xe]6s1 tt 87 Fr [Rn]7s 1 Thus it can be seen that the properties of the Periodic Table. More about these elements an element have periodic dependence upon its will be discussed later. During the description no atomic number and not on relative atomic of their features certain terminology has been mass. used which has been classified in section 3.7. 3.6 ELECTRONIC CONFIGURATIONS AND TYPES OF ELEMENTS: 3.6.1 The s-Block Elements s-, p-, d-, f- BLOCKS The aufbau (build up) principle and the The elements of Group 1 (alkali metals) and electronic configuration of atoms provide a Group 2 (alkaline earth metals) which have ns1  80 h ed pu T is re ER bl be C o N © CHEMISTRY tt Fig. 3.3 The types of elements in the Periodic Table based on the orbitals that are being filled. Also shown is the broad division of elements into METALS ( ) , NON-METALS ( ) and METALLOIDS ( ). no CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 81 and ns2 outermost electronic configuration used as catalysts. However, Zn, Cd and Hg belong to the s-Block Elements. They are all which have the electronic configuration, reactive metals with low ionization enthalpies. (n-1) d10ns2 do not show most of the properties They lose the outermost electron(s) readily to of transition elements. In a way, transition form 1+ ion (in the case of alkali metals) or 2+ metals form a bridge between the chemically ion (in the case of alkaline earth metals). The active metals of s-block elements and the less metallic character and the reactivity increase active elements of Groups 13 and 14 and thus as we go down the group. Because of high take their familiar name “ Transition d reactivity they are never found pure in nature. Elements”. The compounds of the s-block elements, with 3.6.4 The f-Block Elements he the exception of those of lithium and beryllium (Inner-Transition Elements) are predominantly ionic. The two rows of elements at the bottom of the 3.6.2 The p-Block Elements Periodic Table, called the Lanthanoids, Ce(Z = 58) – Lu(Z = 71) and Actinoids, pu T The p -Block Elements comprise those Th(Z = 90) – Lr (Z = 103) are characterised by is belonging to Group 13 to 18 and these the outer electronic configuration (n-2)f 1-14 together with the s-Block Elements are called re ER the Representative Elements or Main Group (n-1)d0–1 ns2. The last electron added to each element is filled in f- orbital. These two series bl Elements. The outermost electronic of elements are hence called the Inner- configuration varies from ns 2np1 to ns2np6 in Transition Elements (f-Block Elements). each period. At the end of each period is a noble They are all metals. Within each series, the gas element with a closed valence shell ns2np6 properties of the elements are quite similar. The be C configuration. All the orbitals in the valence chemistry of the early actinoids is more shell of the noble gases are completely filled complicated than the corresponding by electrons and it is very difficult to alter this lanthanoids, due to the large number of o N stable arrangement by the addition or removal oxidation states possible for these actinoid of electrons. The noble gases thus exhibit very elements. Actinoid elements are radioactive. low chemical reactivity. Preceding the noble gas Many of the actinoid elements have been made family are two chemically important groups of only in nanogram quantities or even less by non-metals. They are the halogens (Group 17) © nuclear reactions and their chemistry is not and the chalcogens (Group 16). These two fully studied. The elements after uranium are groups of elements have highly negative called Transuranium Elements. electron gain enthalpies and readily add one or two electrons respectively to attain the stable Problem 3.3 noble gas configuration. The non-metallic The elements Z = 117 and 120 have not character increases as we move from left to right yet been discovered. In which family / across a period and metallic character increases group would you place these elements as we go down the group. and also give the electronic configuration tt 3.6.3 The d-Block Elements (Transition in each case. Elements) Solution These are the elements of Group 3 to 12 in the We see from Fig. 3.2, that element with Z no centre of the Periodic Table. These are = 117, would belong to the halogen family characterised by the filling of inner d orbitals (Group 17) and the electronic by electrons and are therefore referred to as configuration would be [Rn] d-Block Elements. These elements have the 5f 146d107s 27p5. The element with Z = 120, general outer electronic configuration will be placed in Group 2 (alkaline earth (n-1)d1-10ns0-2. They are all metals. They mostly metals), and will have the electronic form coloured ions, exhibit variable valence configuration [Uuo]8s2. (oxidation states), paramagnetism and oftenly 82 CHEMISTRY 3.6.5 Metals, Non-metals and Metalloids from left to right. Hence the order of increasing metallic character is: P < Si < In addition to displaying the classification of Be < Mg < Na. elements into s-, p-, d-, and f-blocks, Fig. 3.3 shows another broad classification of elements 3.7 PERIODIC TRENDS IN PROPERTIES based on their properties. The elements can OF ELEMENTS be divided into Metals and Non-Metals. Metals comprise more than 78% of all known elements There are many observable patterns in the and appear on the left side of the Periodic physical and chemical properties of elements d Table. Metals are usually solids at room as we descend in a group or move across a temperature [mercury is an exception; gallium period in the Periodic Table. For example, he and caesium also have very low melting points within a period, chemical reactivity tends to be (303K and 302K, respectively)]. Metals usually high in Group 1 metals, lower in elements have high melting and boiling points. They are towards the middle of the table, and increases good conductors of heat and electricity. They to a maximum in the Group 17 non-metals. pu T are malleable (can be flattened into thin sheets Likewise within a group of representative is by hammering) and ductile (can be drawn into metals (say alkali metals) reactivity increases re ER wires). In contrast, non-metals are located at the top right hand side of the Periodic Table. on moving down the group, whereas within a group of non-metals (say halogens), reactivity bl In fact, in a horizontal row, the property of decreases down the group. But why do the elements change from metallic on the left to properties of elements follow these trends? And non-metallic on the right. Non-metals are how can we explain periodicity? To answer usually solids or gases at room temperature these questions, we must look into the theories be C with low melting and boiling points (boron and of atomic structure and properties of the atom. carbon are exceptions). They are poor In this section we shall discuss the periodic conductors of heat and electricity. Most non- trends in certain physical and chemical o N metallic solids are brittle and are neither properties and try to explain them in terms of malleable nor ductile. The elements become number of electrons and energy levels. more metallic as we go down a group; the non- metallic character increases as one goes from 3.7.1 Trends in Physical Properties © left to right across the Periodic Table. The There are numerous physical properties of change from metallic to non-metallic character elements such as melting and boiling points, is not abrupt as shown by the thick zig-zag heats of fusion and vaporization, energy of line in Fig. 3.3. The elements (e.g., silicon, atomization, etc. which show periodic germanium, arsenic, antimony and tellurium) variations. However, we shall discuss the bordering this line and running diagonally periodic trends with respect to atomic and ionic across the Periodic Table show properties that are characteristic of both metals and non- radii, ionization enthalpy, electron gain metals. These elements are called Semi-metals enthalpy and electronegativity. tt or Metalloids. (a) Atomic Radius Problem 3.4 You can very well imagine that finding the size Considering the atomic number and of an atom is a lot more complicated than no position in the periodic table, arrange the measuring the radius of a ball. Do you know following elements in the increasing order why? Firstly, because the size of an atom of metallic character : Si, Be, Mg, Na, P. (~ 1.2 Å i.e., 1.2 × 10–10 m in radius) is very small. Secondly, since the electron cloud Solution surrounding the atom does not have a sharp Metallic character increases down a group boundary, the determination of the atomic size and decreases along a period as we move cannot be precise. In other words, there is no CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 83 practical way by which the size of an individual explain these trends in terms of nuclear charge atom can be measured. However, an estimate and energy level. The atomic size generally of the atomic size can be made by knowing the decreases across a period as illustrated in distance between the atoms in the combined Fig. 3.4(a) for the elements of the second period. state. One practical approach to estimate the It is because within the period the outer size of an atom of a non-metallic element is to electrons are in the same valence shell and the measure the distance between two atoms when effective nuclear charge increases as the atomic they are bound together by a single bond in a number increases resulting in the increased d covalent molecule and from this value, the attraction of electrons to the nucleus. Within a “Covalent Radius” of the element can be family or vertical column of the periodic table, calculated. For example, the bond distance in he the atomic radius increases regularly with the chlorine molecule (Cl2) is 198 pm and half atomic number as illustrated in Fig. 3.4(b). For this distance (99 pm), is taken as the atomic alkali metals and halogens, as we descend the radius of chlorine. For metals, we define the groups, the principal quantum number (n) pu T term “Metallic Radius” which is taken as half increases and the valence electrons are farther is the internuclear distance separating the metal from the nucleus. This happens because the cores in the metallic crystal. For example, the re ER distance between two adjacent copper atoms inner energy levels are filled with electrons, which serve to shield the outer electrons from bl in solid copper is 256 pm; hence the metallic radius of copper is assigned a value of 128 pm. the pull of the nucleus. Consequently the size For simplicity, in this book, we use the term of the atom increases as reflected in the atomic Atomic Radius to refer to both covalent or radii. Note that the atomic radii of noble gases be C metallic radius depending on whether the element is a non-metal or a metal. Atomic radii are not considered here. Being monoatomic, can be measured by X-ray or other their (non-bonded radii) values are very large. o N spectroscopic methods. In fact radii of noble gases should be compared The atomic radii of a few elements are listed not with the covalent radii but with the van der in Table 3.6. Two trends are obvious. We can Waals radii of other elements. Table 3.6(a) Atomic Radii/pm Across the Periods © Atom (Period II) Li Be B C N O F Atomic radius 152 111 88 77 74 66 64 Atom (Period III) Na Mg Al Si P S Cl Atomic radius 186 160 143 117 110 104 99 Table 3.6(b) Atomic Radii/pm Down a Family tt Atom Atomic Atom Atomic (Group I) Radius (Group 17) Radius no Li 152 F 64 Na 186 Cl 99 K 231 Br 114 Rb 244 I 133 Cs 262 At 140 84 CHEMISTRY d he pu T is Fig. 3.4 (a) Variation of atomic radius with Fig. 3.4 (b) Variation of atomic radius with re ER atomic number across the second atomic number for alkali metals bl period and halogens (b) Ionic Radius attraction of the electrons to the nucleus. Anion with the greater negative charge will have the The removal of an electron from an atom results larger radius. In this case, the net repulsion of in the formation of a cation, whereas gain of be C the electrons will outweigh the nuclear charge an electron leads to an anion. The ionic radii and the ion will expand in size. can be estimated by measuring the distances o N between cations and anions in ionic crystals. Problem 3.5 In general, the ionic radii of elements exhibit the same trend as the atomic radii. A cation is Which of the following species will have smaller than its parent atom because it has the largest and the smallest size? fewer electrons while its nuclear charge remains Mg, Mg2+, Al, Al3+. © the same. The size of an anion will be larger Solution than that of the parent atom because the addition of one or more electrons would result Atomic radii decrease across a period. in increased repulsion among the electrons Cations are smaller than their parent and a decrease in effective nuclear charge. For atoms. Among isoelectronic species, the – example, the ionic radius of fluoride ion (F ) is one with the larger positive nuclear charge 136 pm whereas the atomic radius of fluorine will have a smaller radius. is only 64 pm. On the other hand, the atomic Hence the largest species is Mg; the radius of sodium is 186 pm compared to the smallest one is Al3+. tt + ionic radius of 95 pm for Na. When we find some atoms and ions which (c) Ionization Enthalpy contain the same number of electrons, we call no them isoelectronic species*. For example, A quantitative measure of the tendency of an O2–, F–, Na + and Mg2+ have the same number of element to lose electron is given by its electrons (10). Their radii would be different Ionization Enthalpy. It represents the energy because of their different nuclear charges. The required to remove an electron from an isolated cation with the greater positive charge will have gaseous atom (X) in its ground state. In other a smaller radius because of the greater words, the first ionization enthalpy for an * Two or more species with same number of atoms, same number of valence electrons and same structure, regardless of the nature of elements involved. CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 85 element X is the enthalpy change (∆i H) for the reaction depicted in equation 3.1. X(g) → X+(g) + e– (3.1) The ionization enthalpy is expressed in units of kJ mol–1. We can define the second ionization enthalpy as the energy required to remove the second most loosely bound d electron; it is the energy required to carry out the reaction shown in equation 3.2. he + 2+ – X (g) → X (g) + e (3.2) Energy is always required to remove Fig. 3.5 Variation of first ionization enthalpies pu T electrons from an atom and hence ionization (∆i H) with atomic number for elements is enthalpies are always positive. The second with Z = 1 to 60 ionization enthalpy will be higher than the first re ER ionization enthalpy because it is more difficult with their high reactivity. In addition, you will bl to remove an electron from a positively charged notice two trends the first ionization enthalpy ion than from a neutral atom. In the same way generally increases as we go across a period the third ionization enthalpy will be higher than and decreases as we descend in a group. These the second and so on. The term “ionization trends are illustrated in Figs. 3.6(a) and 3.6(b) enthalpy”, if not qualified, is taken as the first respectively for the elements of the second be C ionization enthalpy. period and the first group of the periodic table. The first ionization enthalpies of elements You will appreciate that the ionization enthalpy and atomic radius are closely related o N having atomic numbers up to 60 are plotted in Fig. 3.5. The periodicity of the graph is quite properties. To understand these trends, we striking. You will find maxima at the noble gases have to consider two factors : (i) the attraction which have closed electron shells and very of electrons towards the nucleus, and (ii) the stable electron configurations. On the other repulsion of electrons from each other. The © hand, minima occur at the alkali metals and effective nuclear charge experienced by a their low ionization enthalpies can be correlated valence electron in an atom will be less than tt no 3.6 (a) 3.6 (b) Fig. 3.6(a) First ionization enthalpies (∆ i H) of elements of the second period as a function of atomic number (Z) and Fig. 3.6(b) ∆i H of alkali metals as a function of Z. 86 CHEMISTRY the actual charge on the nucleus because of to remove the 2p-electron from boron compared “shielding” or “screening” of the valence to the removal of a 2s- electron from beryllium. electron from the nucleus by the intervening Thus, boron has a smaller first ionization core electrons. For example, the 2s electron in enthalpy than beryllium. Another “anomaly” lithium is shielded from the nucleus by the is the smaller first ionization enthalpy of oxygen inner core of 1s electrons. As a result, the compared to nitrogen. This arises because in valence electron experiences a net positive the nitrogen atom, three 2p-electrons reside in charge which is less than the actual charge of different atomic orbitals (Hund’s rule) whereas d +3. In general, shielding is effective when the in the oxygen atom, two of the four 2p-electrons orbitals in the inner shells are completely filled. must occupy the same 2p-orbital resulting in an increased electron-electron repulsion. he This situation occurs in the case of alkali metals which have single outermost ns-electron Consequently, it is easier to remove the fourth preceded by a noble gas electronic 2p-electron from oxygen than it is, to remove configuration. one of the three 2p-electrons from nitrogen. pu T When we move from lithium to fluorine is Problem 3.6 across the second period, successive electrons The first ionization enthalpy (∆i H ) values re ER are added to orbitals in the same principal quantum level and the shielding of the nuclear of the third period elements, Na, Mg and bl Si are respectively 496, 737 and 786 kJ charge by the inner core of electrons does not mol–1. Predict whether the first ∆ i H value increase very much to compensate for the increased attraction of the electron to the for Al will be more close to 575 or 760 kJ mol–1 ? Justify your answer. nucleus. Thus, across a period, increasing be C nuclear charge outweighs the shielding. Solution Consequently, the outermost electrons are held –1 It will be more close to 575 kJ mol. The more and more tightly and the ionization value for Al should be lower than that of o N enthalpy increases across a period. As we go Mg because of effective shielding of 3p down a group, the outermost electron being electrons from the nucleus by increasingly farther from the nucleus, there is 3s-electrons. an increased shielding of the nuclear charge © by the electrons in the inner levels. In this case, (d) Electron Gain Enthalpy increase in shielding outweighs the increasing When an electron is added to a neutral gaseous nuclear charge and the removal of the atom (X) to convert it into a negative ion, the outermost electron requires less energy down enthalpy change accompanying the process is a group. defined as the Electron Gain Enthalpy (∆ egH). From Fig. 3.6(a), you will also notice that Electron gain enthalpy provides a measure of the first ionization enthalpy of boron (Z = 5) is the ease with which an atom adds an electron slightly less than that of beryllium (Z = 4) even to form anion as represented by equation 3.3. though the former has a greater nuclear charge. X(g) + e – → X –(g) tt When we consider the same principal quantum (3.3) level, an s-electron is attracted to the nucleus Depending on the element, the process of more than a p-electron. In beryllium, the adding an electron to the atom can be either electron removed during the ionization is an endothermic or exothermic. For many elements no s-electron whereas the electron removed during energy is released when an electron is added ionization of boron is a p-electron. The to the atom and the electron gain enthalpy is penetration of a 2s-electron to the nucleus is negative. For example, group 17 elements (the more than that of a 2p-electron; hence the 2p halogens) have very high negative electron gain electron of boron is more shielded from the enthalpies because they can attain stable noble nucleus by the inner core of electrons than the gas electronic configurations by picking up an 2s electrons of beryllium. Therefore, it is easier electron. On the other hand, noble gases have CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 87 Table 3.7 Electron Gain Enthalpies* / (kJ mol–1 ) of Some Main Group Elements Group 1 ∆ eg H Group 16 ∆ eg H Group 17 ∆ eg H Group 0 ∆ eg H H – 73 He + 48 Li – 60 O – 141 F – 328 Ne + 116 Na – 53 S – 200 Cl – 349 Ar + 96 K – 48 Se – 195 Br – 325 Kr + 96 d Rb – 47 Te – 190 I – 295 Xe + 77 he Cs – 46 Po – 174 At – 270 Rn + 68 large positive electron gain enthalpies because Problem 3.7 the electron has to enter the next higher pu T principal quantum level leading to a very Which of the following will have the most is unstable electronic configuration. It may be negative electron gain enthalpy and which the least negative? re ER noted that electron gain enthalpies have large negative values toward the upper right of the P, S, Cl, F. bl periodic table preceding the noble gases. Explain your answer. The variation in electron gain enthalpies of elements is less systematic than for ionization Solution enthalpies. As a general rule, electron gain Electron gain enthalpy generally becomes be C enthalpy becomes more negative with increase more negative across a period as we move in the atomic number across a period. The from left to right. Within a group, electron effective nuclear charge increases from left to gain enthalpy becomes less negative down o N right across a period and consequently it will a group. However, adding an electron to be easier to add an electron to a smaller atom the 2p-orbital leads to greater repulsion since the added electron on an average would than adding an electron to the larger be closer to the positively charged nucleus. We 3p-orbital. Hence the element with most should also expect electron gain enthalpy to © negative electron gain enthalpy is chlorine; become less negative as we go down a group the one with the least negative electron because the size of the atom increases and the gain enthalpy is phosphorus. added electron would be farther from the nucleus. This is generally the case (Table 3.7). (e) Electronegativity However, electron gain enthalpy of O or F is less negative than that of the succeeding A qualitative measure of the ability of an atom element. This is because when an electron is in a chemical compound to attract shared added to O or F, the added electron goes to the electrons to itself is called electronegativity. smaller n = 2 quantum level and suffers Unlike ionization enthalpy and electron gain tt significant repulsion from the other electrons enthalpy, it is not a measureable quantity. present in this level. For the n = 3 quantum However, a number of numerical scales of level (S or Cl), the added electron occupies a electronegativity of elements viz., Pauling scale, larger region of space and the electron-electron Mulliken-Jaffe scale, Allred-Rochow scale have no repulsion is much less. been developed. The one which is the most * In many books, the negative of the enthalpy change for the process depicted in equation 3.3 is defined as the ELECTRON AFFINITY (Ae ) of the atom under consideration. If energy is released when an electron is added to an atom, the electron affinity is taken as positive, contrary to thermodynamic convention. If energy has to be supplied to add an electron to an atom, then the electron affinity of the atom is assigned a negative sign. However, electron affinity is defined as absolute zero and, therefore at any other temperature (T) heat capacities of the reactants and the products have to be taken into account in ∆egH = –Ae – 5/2 RT. 88 CHEMISTRY widely used is the Pauling scale. Linus Pauling, electrons and the nucleus increases as the an American scientist, in 1922 assigned atomic radius decreases in a period. The arbitrarily a value of 4.0 to fluorine, the element electronegativity also increases. On the same considered to have the greatest ability to attract account electronegativity values decrease with electrons. Approximate values for the the increase in atomic radii down a group. The electronegativity of a few elements are given in trend is similar to that of ionization enthalpy. Table 3.8(a) Knowing the relationship between The electronegativity of any given element electronegativity and atomic radius, can you d is not constant; it varies depending on the now visualise the relationship between element to which it is bound. Though it is not electronegativity and non-metallic properties? he a measurable quantity, it does provide a means of predicting the nature of force that holds a pair of atoms together – a relationship that you will pu T explore later. is Electronegativity generally re ER increases across a period from left to right (say from lithium to bl fluorine) and decrease down a group (say from fluorine to astatine) in the periodic table. How can these trends be explained? Can the be C electronegativity be related to atomic radii, which tend to decrease across each period from o N left to right, but increase down each group ? The attraction between the outer (or valence) Fig. 3.7 The periodic trends of elements in the periodic table © Table 3.8(a) Electronegativity Values (on Pauling scale) Across the Periods Atom (Period II) Li Be B C N O F Electronegativity 1.0 1.5 2.0 2.5 3.0 3.5 4.0 Atom (Period III) Na Mg Al Si P S Cl Electronegativity 0.9 1.2 1.5 1.8 2.1 2.5 3.0 Table 3.8(b) Electronegativity Values (on Pauling scale) Down a Family tt Atom Electronegativity Atom Electronegativity (Group I) Value (Group 17) Value no Li 1.0 F 4.0 Na 0.9 Cl 3.0 K 0.8 Br 2.8 Rb 0.8 I 2.5 Cs 0.7 At 2.2 CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 89 Non-metallic elements have strong tendency electronic configuration 2s 22p5, shares one to gain electrons. Therefore, electronegativity electron with oxygen in the OF2 molecule. Being is directly related to that non-metallic highest electronegative element, fluorine is properties of elements. It can be further given oxidation state –1. Since there are two extended to say that the electronegativity is fluorine atoms in this molecule, oxygen with 2 inversely related to the metallic properties of outer electronic configuration 2s 2p4 shares elements. Thus, the increase in two electrons with fluorine atoms and thereby electronegativities across a period is exhibits oxidation state +2. In Na2O, oxygen d accompanied by an increase in non-metallic being more electronegative accepts two properties (or decrease in metallic properties) electrons, one from each of the two sodium of elements. Similarly, the decrease in atoms and, thus, shows oxidation state –2. On he electronegativity down a group is accompanied the other hand sodium with electronic by a decrease in non-metallic properties (or configuration 3s 1 loses one electron to oxygen increase in metallic properties) of elements. and is given oxidation state +1. Thus, the pu T All these periodic trends are summarised oxidation state of an element in a particular is in figure 3.7. compound can be defined as the charge acquired by its atom on the basis of re ER 3.7.2 Periodic Trends in Chemical electronegative consideration from other atoms bl Properties in the molecule. Most of the trends in chemical properties of elements, such as diagonal relationships, inert Problem 3.8 pair effect, effects of lanthanoid contraction etc. Using the Periodic Table, predict the be C will be dealt with along the discussion of each formulas of compounds which might be group in later units. In this section we shall formed by the following pairs of elements; study the periodicity of the valence state shown (a) silicon and bromine (b) aluminium and o N by elements and the anomalous properties of sulphur. the second period elements (from lithium to fluorine). Solution (a) Periodicity of Valence or Oxidation (a) Silicon is group 14 element with a valence of 4; bromine belongs to the © States halogen family with a valence of 1. The valence is the most characteristic property Hence the formula of the compound of the elements and can be understood in terms formed would be SiBr4. of their electronic configurations. The valence of representative elements is usually (though (b) Aluminium belongs to group 13 with not necessarily) equal to the number of a valence of 3; sulphur belongs to electrons in the outermost orbitals and / or group 16 elements with a valence of equal to eight minus the number of outermost 2. Hence, the formula of the compound electrons as shown below. formed would be Al2S3. Nowadays the term oxidation state is tt Some periodic trends observed in the frequently used for valence. Consider the two valence of elements (hydrides and oxides) are oxygen containing compounds: OF2 and Na2O. shown in Table 3.9. Other such periodic trends The order of electronegativity of the three which occur in the chemical behaviour of the no elements involved in these compounds is F > elements are discussed elsewhere in this book. O > Na. Each of the atoms of fluorine, with outer Group 1 2 13 14 15 16 17 18 Number of valence 1 2 3 4 5 6 7 8 electron Valence 1 2 3 4 3,5 2,6 1,7 0,8 90 CHEMISTRY Table 3.9 Periodic Trends in Valence of Elements as shown by the Formulas of Their Compounds Group 1 2 13 14 15 16 17 Formula LiH B2 H6 CH4 NH3 H2O HF of hydride NaH CaH2 AlH3 SiH4 PH3 H2 S HCl KH GeH4 AsH3 H2 Se HBr d SnH4 SbH3 H2 Te HI Formula Li2O MgO B2 O3 CO2 N2O 3, N2 O5 – he of oxide Na2 O CaO Al2O3 SiO2 P4 O6, P4O 10 SO3 Cl2 O 7 K2O SrO Ga2 O3 GeO2 As2O 3, As2 O5 SeO3 – pu T BaO In2O3 SnO2 Sb2O 3, Sb2O 5 TeO3 – is PbO2 Bi2 O3 – – re ER bl There are many elements which exhibit variable following group i.e., magnesium and valence. This is particularly characteristic of aluminium, respectively. This sort of similarity transition elements and actinoids, which we is commonly referred to as diagonal shall study later. relationship in the periodic properties. be C (b) Anomalous Properties of Second Period What are the reasons for the different Elements chemical behaviour of the first member of a group of elements in the s- and p-blocks The first element of each of the groups 1 o N compared to that of the subsequent members (lithium) and 2 (beryllium) and groups 13-17 in the same group? The anomalous behaviour (boron to fluorine) differs in many respects from is attributed to their small size, large charge/ the other members of their respective group. radius ratio and high electronegativity of the For example, lithium unlike other alkali metals, elements. In addition, the first member of © and beryllium unlike other alkaline earth group has only four valence orbitals (2s and metals, form compounds with pronounced 2p) available for bonding, whereas the second covalent character; the other members of these member of the groups have nine valence groups predominantly form ionic compounds. orbitals (3s, 3p, 3d). As a consequence of this, In fact the behaviour of lithium and beryllium the maximum covalency of the first member of is more similar with the second element of the each group is 4 (e.g., boron can only form − Property Element [BF4 ] , whereas the other members of the groups can expand their Metallic radius M/ pm Li Be B valence shell to accommodate more tt 152 111 88 than four pairs of electrons e.g., 3− Na Mg Al aluminium forms [ AlF6 ] ). Furthermore, the first member of no 186 160 143 p-block elements displays greater Li Be ability to form pπ – p π multiple bonds + to itself (e.g., C = C, C ≡ C, N = N, Ionic radius M / pm 76 31 N ≡ Ν) and to other second period Na Mg elements (e.g., C = O, C = N, C ≡ N, N = O) compared to subsequent 102 72 members of the same group. CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 91 Problem 3.9 non-metallic character increases while moving from left to right across the period. The Are the oxidation state and covalency of chemical reactivity of an element can be best 2+ Al in [AlCl(H2 O)5] same ? shown by its reactions with oxygen and Solution halogens. Here, we shall consider the reaction of the elements with oxygen only. Elements on No. The oxidation state of Al is +3 and the two extremes of a period easily combine with covalency is 6. oxygen to form oxides. The normal oxide d formed by the element on extreme left is the 3.7.3 Periodic Trends and Chemical most basic (e.g., Na2O), whereas that formed Reactivity by the element on extreme right is the most he acidic (e.g., Cl2O 7). Oxides of elements in the We have observed the periodic trends in certain centre are amphoteric (e.g., Al2O 3, As2O3) or fundamental properties such as atomic and neutral (e.g., CO, NO, N2O). Amphoteric oxides ionic radii, ionization enthalpy, electron gain pu T behave as acidic with bases and as basic with enthalpy and valence. We know by now that is acids, whereas neutral oxides have no acidic the periodicity is related to electronic or basic properties. re ER configuration. That is, all chemical and physical properties are a manifestation of the bl Problem 3.10 electronic configuration of elements. We shall now try to explore relationships between these Show by a chemical reaction with water fundamental properties of elements with their that Na2O is a basic oxide and Cl2O7 is an chemical reactivity. acidic oxide. be C The atomic and ionic radii, as we know, Solution generally decrease in a period from left to right. Na2O with water forms a strong base As a consequence, the ionization enthalpies o N whereas Cl2O7 forms strong acid. generally increase (with some exceptions as outlined in section 3.7.1(a)) and electron gain Na2O + H2O → 2NaOH enthalpies become more negative across a period. In other words, the ionization enthalpy Cl2O 7 + H2O → 2HClO4 © of the extreme left element in a period is the least and the electron gain enthalpy of the Their basic or acidic nature can be element on the extreme right is the highest qualitatively tested with litmus paper. negative (note : noble gases having completely filled shells have rather positive electron gain Among transition metals (3d series), the change enthalpy values). This results into high in atomic radii is much smaller as compared chemical reactivity at the two extremes and the to those of representative elements across the lowest in the centre. Thus, the maximum period. The change in atomic radii is still chemical reactivity at the extreme left (among smaller among inner-transition metals (4f series). The ionization enthalpies are tt alkali metals) is exhibited by the loss of an electron leading to the formation of a cation intermediate between those of s- and p-blocks. and at the extreme right (among halogens) As a consequence, they are less electropositive shown by the gain of an electron forming an than group 1 and 2 metals. no anion. This property can be related with the In a group, the increase in atomic and ionic reducing and oxidizing behaviour of the radii with increase in atomic number generally elements which you will learn later. However, results in a gradual decrease in ionization here it can be directly related to the metallic enthalpies and a regular decrease (with and non-metallic character of elements. Thus, exception in some third period elements as the metallic character of an element, which is shown in section 3.7.1(d)) in electron gain highest at the extremely left decreases and the enthalpies in the case of main group elements. 92 CHEMISTRY Thus, the metallic character increases down will learn later. In the case of transition the group and non-metallic character elements, however, a reverse trend is observed. decreases. This trend can be related with their This can be explained in terms of atomic size reducing and oxidizing property which you and ionization enthalpy. SUMMARY d In this Unit, you have studied the development of the Periodic Law and the Periodic he Table. Mendeleev’s Periodic Table was based on atomic masses. Modern Periodic Table arranges the elements in the order of their atomic numbers in seven horizontal rows (periods) and eighteen vertical columns (groups or families). Atomic numbers in a period are consecutive, whereas in a group they increase in a pattern. Elements of the same group have similar valence shell electronic configuration and, therefore, exhibit similar pu T chemical properties. However, the elements of the same period have incrementally is increasing number of electrons from left to right, and, therefore, have different valencies. re ER Four types of elements can be recognized in the periodic table on the basis of their electronic configurations. These are s-block, p-block, d-block and f -block elements. bl Hydrogen with one electron in the 1s orbital occupies a unique position in the periodic table. Metals comprise more than seventy eight per cent of the known elements. Non- metals, which are located at the top of the periodic table, are less than twenty in number. Elements which lie at the border line between metals and non-metals (e.g., Si, Ge, As) are called metalloids or semi-metals. Metallic character increases with increasing atomic be C number in a group whereas decreases from left to right in a period. The physical and chemical properties of elements vary periodically with their atomic numbers. Periodic trends are observed in atomic sizes, ionization enthalpies, electron o N gain enthalpies, electronegativity and valence. The atomic radii decrease while going from left to right in a period and increase with atomic number in a group. Ionization enthalpies generally increase across a period and decrease down a group. Electronegativity also shows a similar trend. Electron gain enthalpies, in general, become more negative © across a period and less negative down a group. There is some periodicity in valence, for example, among representative elements, the valence is either equal to the number of electrons in the outermost orbitals or eight minus this number. Chemical reactivity is hightest at the two extremes of a period and is lowest in the centre. The reactivity on the left extreme of a period is because of the ease of electron loss (or low ionization enthalpy). Highly reactive elements do not occur in nature in free state; they usually occur in the combined form. Oxides formed of the elements on the left are basic and of the elements on the right are acidic in nature. Oxides of elements in the centre are amphoteric or neutral. tt EXERCISES no 3.1 What is the basic theme of organisation in the periodic table? 3.2 Which important property did Mendeleev use to classify the elements in his periodic table and did he stick to that? 3.3 What is the basic difference in approach between the Mendeleev’s Periodic Law and the Modern Periodic Law? 3.4 On the basis of quantum numbers, justify that the sixth period of the periodic table should have 32 elements. CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 93 3.5 In terms of period and group where would you locate the element with Z =114? 3.6 Write the atomic number of the element present in the third period and seventeenth group of the periodic table. 3.7 Which element do you think would have been named by (i) Lawrence Berkeley Laboratory (ii) Seaborg’s group? 3.8 Why do elements in the same group have similar physical and chemical properties? 3.9 What does atomic radius and ionic radius really mean to you? d 3.10 How do atomic radius vary in a period and in a group? How do you explain the variation? he 3.11 What do you understand by isoelectronic species? Name a species that will be isoelectronic with each of the following atoms or ions. – + (i) F (ii) Ar (iii) Mg2+ (iv) Rb 3.12 Consider the following species : pu T 3– 2– – + 2+ 3+ N , O , F , Na , Mg and Al is (a) What is common in them? re ER (b) Arrange them in the order of increasing ionic radii. bl 3.13 Explain why cation are smaller and anions larger in radii than their parent atoms? 3.14 What is the significance of the terms — ‘isolated gaseous atom’ and ‘ground state’ while defining the ionization enthalpy and electron gain enthalpy? Hint : Requirements for comparison purposes. be C 3.15 Energy of an electron in the ground state of the hydrogen atom is –2.18×10–18J. Calculate the ionization enthalpy of atomic hydrogen in terms of J mol–1. Hint: Apply the idea of mole concept to derive the answer. o N 3.16 Among the second period elements the actual ionization enthalpies are in the order Li < B < Be < C < O < N < F < Ne. Explain why (i) Be has higher ∆i H than B © (ii) O has lower ∆i H than N and F? 3.17 How would you explain the fact that the first ionization enthalpy of sodium is lower than that of magnesium but its second ionization enthalpy is higher than that of magnesium? 3.18 What are the various factors due to which the ionization enthalpy of the main group elements tends to decrease down a group? 3.19 The first ionization enthalpy values (in kJ mol–1) of group 13 elements are : B Al Ga In Tl 801 577 579 558 589 tt How would you explain this deviation from the general trend ? 3.20 Which of the following pairs of elements would have a more negative electron gain enthalpy? no (i) O or F (ii) F or Cl 3.21 Would you expect the second electron gain enthalpy of O as positive, more negative or less negative than the first? Justify your answer. 3.22 What is the basic difference between the terms electron gain enthalpy and electronegativity? 3.23 How would you react to the statement that the electronegativity of N on Pauling scale is 3.0 in all the nitrogen compounds? 94 CHEMISTRY 3.24 Describe the theory associated with the radius of an atom as it (a) gains an electron (b) loses an electron 3.25 Would you expect the first ionization enthalpies for two isotopes of the same element to be the same or differ ent? Justify your answer. 3.26 What are the major differences between metals and non-metals? 3.27 Use the periodic table to answer the following questions. (a) Identify an element with five electrons in the outer subshell. d (b) Identify an element that would tend to lose two electrons. (c) Identify an element that would tend to gain two electrons. he (d) Identify the group having metal, non-metal, liquid as well as gas at the room temperature. 3.28 The increasing order of reactivity among group 1 elements is Li < Na < K < Rb CI > Br > I. Explain. pu T 3.29 Write the general outer electronic configuration of s-, p-, d- and f- block elements. is 3.30 Assign the position of the element having outer electronic configuration re ER (i) ns2 np4 for n=3 (ii) (n-1)d2 ns2 for n=4, and (iii) (n-2) f 7 (n-1)d1 ns2 for n=6, in the periodic table. bl 3.31 The first (∆i H 1) and the second (∆i H2) ionization enthalpies (in kJ mol –1) and the (∆egH) electron gain enthalpy (in kJ mol–1) of a few elements are given below: Elements ∆H 1 ∆H 2 ∆ egH I 520 7300 –60 be C II 419 3051 –48 III 1681 3374 –328 IV 1008 1846 –295 o N V 2372 5251 +48 VI 738 1451 –40 Which of the above elements is likely to be : (a) the least reactive element. © (b) the most reactive metal. (c) the most reactive non-metal. (d) the least reactive non-metal. (e) the metal which can form a stable binary halide of the formula MX 2(X=halogen). (f) the metal which can form a predominantly stable covalent halide of the formula MX (X=halogen)? 3.32 Predict the formulas of the stable binary compounds that would be formed by the combination of the following pairs of elements. (a) Lithium and oxygen (b) Magnesium and nitrogen tt (c) Aluminium and iodine (d) Silicon and oxygen (e) Phosphorus and fluorine (f) Element 71 and fluorine 3.33 In the modern periodic table, the period indicates the value of : no (a) atomic number (b) atomic mass (c) principal quantum number (d) azimuthal quantum number. 3.34 Which of the following statements related to the modern periodic table is incorrect? (a) The p-block has 6 columns, because a maximum of 6 electrons can occupy all the orbitals in a p-shell. CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPER TIES 95 (b) The d-block has 8 columns, because a maximum of 8 electrons can occupy all the orbitals in a d-subshell. (c) Each block contains a number of columns equal to the number of electrons that can occupy that subshell. (d) The block indicates value of azimuthal quantum number (l) for the last subshell that received electrons in building up the electronic configuration. 3.35 Anything that influences the valence electrons will affect the chemistry d of the element. Which one of the following factors does not affect the valence shell? (a) Valence principal quantum number (n) he (b) Nuclear charge (Z ) (c) Nuclear mass (d) Number of core electrons. pu T – 3.36 The size of isoelectronic species — F , Ne and Na+ is affected by is (a) nuclear charge ( Z ) (b) valence principal quantum number (n) re ER (c) electron-electron interaction in the outer orbitals bl (d) none of the factors because their size is the same. 3.37 Which one of the following statements is incorrect in relation to ionization enthalpy? (a) Ionization enthalpy increases for each successive electron. be C (b) The greatest increase in ionization enthalpy is experienced on removal of electron from core noble gas configuration. o N (c) End of valence electrons is marked by a big jump in ionization enthalpy. (d) Removal of electron from orbitals bearing lower n value is easier than from orbital having higher n value. ©