Unit 3 Classification of Elements and Periodicity in Properties PDF

Unit 3 Classification of Elements and Periodicity in Properties The Periodic Table is arguably the m...

Unit 3 Classification of Elements and Periodicity in Properties The Periodic Table is arguably the most important concept in chemistry, both in principle and in practice. It is the everyday support for students, it suggests new avenues of research to professionals, and it provides a succinct After studying this Unit, you will be organization of the whole of chemistry. It is a remarkable able to demonstration of the fact that the chemical elements are not a random cluster of entities but instead display trends appreciate how the concept of and lie together in families. An awareness of the Periodic grouping elements in accordance Table is essential to anyone who wishes to disentangle to their properties led to the the world and see how it is built up from the fundamental development of Periodic Table. building blocks of the chemistry, the chemical elements. understand the Periodic Law; understand the significance of Glenn T. Seaborg atomic number and electronic configuration as the basis for periodic classification; In this Unit, we will study the historical development of the name the elements with 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 follows as a logical consequence of the electronic classify elements into s, p, d, f 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 We know by now that the elements are the basic units of and correlate it with their all types of matter. In 1800, only 31 elements were known. occurrence in nature; By  1865, the number of identified elements had more than explain the relationship between doubled to 63. At present 114 elements are known. Of ionization enthalpy and metallic them, the recently discovered elements are man-made. character; Efforts to synthesise new elements are continuing. With use scientific vocabulary such a large number of elements it is very difficult to appropriately to communicate study individually the chemistry of all these elements and ideas related to certain important their innumerable compounds individually. To ease out properties of atoms e.g., this problem, scientists searched for a systematic way to atomic/ionic radii, ionization organise their knowledge by classifying the elements. Not enthalpy, electron gain enthalpy, electronegativity, valence of only that it would rationalize known chemical facts about elements. elements, but even predict new ones for undertaking further study. 2024-25 Unit 3.indd 74 9/9/2022 4:36:12 PM  Classification of Elements and Periodicity in Properties 75 3.2 GENESIS OF PERIODIC the periodic recurrence of properties. This CLASSIFICATION also did not attract much attention. The Classification of elements into groups and English chemist, John Alexander Newlands development of Periodic Law and Periodic in 1865 profounded the Law of Octaves. He Table are the consequences of systematising arranged the elements in increasing order the knowledge gained by a number of of their atomic weights and noted that every scientists through their observations and eighth element had properties similar to the experiments. The German chemist, Johann first element (Table 3.2). The relationship was Dobereiner in early 1800’s was the first to just like every eighth note that resembles the consider the idea of trends among properties first in octaves of music. Newlands’s Law of of elements. By 1829 he noted a similarity Octaves seemed to be true only for elements among the physical and chemical properties up to calcium. Although his idea was not of several groups of three elements (Triads). In widely accepted at that time, he, for his work, each case, he noticed that the middle element was later awarded Davy Medal in 1887 by the of each of the Triads had an atomic weight Royal Society, London. about half way between the atomic weights of The Periodic Law, as we know it today the other two (Table 3.1). Also the properties owes its development to the Russian chemist, of the middle element were in between those Dmitri Mendeleev (1834-1907) and the of the other two members. Since Dobereiner’s German chemist, Lothar Meyer (1830-1895). Table 3.1 Dobereiner’s Triads Atomic Atomic Atomic Element Element Element weight weight weight Li 7 Ca 40 Cl 35.5 Na 23 Sr 88 Br 80 K 39 Ba 137 I 127 relationship, referred to as the Law of Triads, Working independently, both the chemists in seemed to work only for a few elements, it was 1869 proposed that on arranging elements in dismissed as coincidence. The next reported the increasing order of their atomic weights, attempt to classify elements was made by a similarities appear in physical and chemical French geologist, A.E.B. de Chancourtois in properties at regular intervals. Lothar Meyer 1862. He arranged the then known elements plotted the physical properties such as in order of increasing atomic weights and atomic volume, melting point and boiling made a cylindrical table of elements to display point against atomic weight and obtained Table 3.2 Newlands’ Octaves Element Li Be B C N O F 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 2024-25 Unit 3.indd 75 9/9/2022 4:36:12 PM  76 chemistry a periodically repeated pattern. Unlike classification if the order of atomic weight Newlands, Lothar Meyer observed a change was strictly followed. He ignored the order in length of that repeating pattern. By 1868, of atomic weights, thinking that the atomic Lothar Meyer had developed a table of the measurements might be incorrect, and placed elements that closely resembles the Modern the elements with similar properties together. Periodic Table. However, his work was not For example, iodine with lower atomic weight published until after the work of Dmitri than that of tellurium (Group VI) was placed Mendeleev, the scientist who is generally in Group VII along with fluorine, chlorine, credited with the development of the Modern bromine because of similarities in properties Periodic Table. (Fig. 3.1). At the same time, keeping his While Dobereiner initiated the study of primary aim of arranging the elements of periodic relationship, it was Mendeleev who similar properties in the same group, he was responsible for publishing the Periodic proposed that some of the elements were Law for the first time. It states as follows : still undiscovered and, therefore, left several gaps in the table. For example, both gallium The properties of the elements are and germanium were unknown at the time a periodic function of their atomic Mendeleev published his Periodic Table. weights. He left the gap under aluminium and a gap Mendeleev arranged elements in horizontal under silicon, and called these elements rows and vertical columns of a table in order Eka-Aluminium and Eka-Silicon. Mendeleev of their increasing atomic weights in such a predicted not only the existence of gallium and way that the elements with similar properties germanium, but also described some of their occupied the same vertical column or group. general physical properties. These elements Mendeleev’s system of classifying elements were discovered later. Some of the properties was more elaborate than that of Lothar predicted by Mendeleev for these elements Meyer’s. He fully recognized the significance and those found experimentally are listed in of periodicity and used broader range of Table 3.3. physical and chemical properties to classify the elements. In particular, Mendeleev relied The boldness of Mendeleev’s quantitative on the similarities in the empirical formulas predictions and their eventual success and properties of the compounds formed by made him and his Periodic Table famous. the elements. He realized that some of the Mendeleev’s Periodic Table published in 1905 elements did not fit in with his scheme of is shown in Fig. 3.1. Table 3.3 Mendeleev’s Predictions for the Elements Eka-aluminium (Gallium) and Eka-silicon (Germanium) Eka-aluminium Gallium Eka-silicon Germanium Property (predicted) (found) (predicted) (found) Atomic weight 68 70 72 72.6 Density/(g/cm3) 5.9 5.94 5.5 5.36 Melting point/K Low 302.93 High 1231 Formula of oxide E2O3 Ga2O3 EO2 GeO2 Formula of chloride E Cl3 GaCl3 ECl4 GeCl4 2024-25 Unit 3.indd 76 9/9/2022 4:36:12 PM Unit 3.indd 77 PERIODIC SYSTEM OF THE ELEMENTS IN GROUPS AND SERIES 2024-25 Classification of Elements and Periodicity in Properties Fig. 3.1 Mendeleev’s Periodic Table published earlier 77 9/9/2022 4:36:12 PM  78 chemistry 3.3 MODERN PERIODIC LAW AND THE physical and chemical properties of elements PRESENT FORM OF THE PERIODIC and their compounds. TABLE Numerous forms of Periodic Table have We must bear in mind that when Mendeleev been devised from time to time. Some developed his Periodic Table, chemists forms emphasise chemical reactions and knew nothing about the internal structure valence, whereas others stress the electronic of atom. However, the beginning of the 20th configuration of elements. A modern version, century witnessed profound developments the so-called “long form” of the Periodic in theories about sub-atomic particles. In Table of the elements (Fig. 3.2), is the most 1913, the English physicist, Henry Moseley convenient and widely used. The horizontal observed regularities in the characteristic rows (which Mendeleev called series) are X-ray spectra of the elements. A plot of called periods and the vertical columns, (where is frequency of X-rays emitted) groups. Elements having similar outer against atomic number (Z) gave a straight electronic configurations in their atoms line and not the plot of vs atomic mass. are arranged in vertical columns, referred He thereby showed that the atomic number to as groups or families. According to the is a more fundamental property of an element recommendation of International Union of Pure and Applied Chemistry (IUPAC), the than its atomic mass. Mendeleev’s Periodic groups are numbered from 1 to 18 replacing Law was, therefore, accordingly modified. This the older notation of groups IA … VIIA, VIII, is known as the Modern Periodic Law and IB … VIIB and 0. can be stated as : There are altogether seven periods. The The physical and chemical properties period number corresponds to the highest of the elements are periodic functions principal quantum number (n) of the elements of their atomic numbers. in the period. The first period contains 2 The Periodic Law revealed important elements. The subsequent periods consists of analogies among the 94 naturally occurring 8, 8, 18, 18 and 32 elements, respectively. The elements (neptunium and plutonium like seventh period is incomplete and like the sixth actinium and protoactinium are also found period would have a theoretical maximum in pitch blende – an ore of uranium). It (on the basis of quantum numbers) of 32 stimulated renewed interest in Inorganic elements. In this form of the Periodic Table, Chemistry and has carried into the present 14 elements of both sixth and seventh periods with the creation of artificially produced (lanthanoids and actinoids, respectively) are short-lived elements. placed in separate panels at the bottom*. You may recall that the atomic number 3.4 NOMENCLATURE OF ELEMENTS is equal to the nuclear charge (i.e., number WITH ATOMIC NUMBERS > 100 of protons) or the number of electrons in a neutral atom. It is then easy to visualize The naming of the new elements had been the significance of quantum numbers and traditionally the privilege of the discoverer electronic configurations in periodicity of (or discoverers) and the suggested name was elements. In fact, it is now recognized that the ratified by the IUPAC. In recent years this has Periodic Law is essentially the consequence led to some controversy. The new elements of the periodic variation in electronic with very high atomic numbers are so unstable configurations, which indeed determine the that only minute quantities, sometimes only * Glenn T. Seaborg’s work in the middle of the 20th 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 Seaborgium (Sg) in his honour. 2024-25 Unit 3.indd 78 9/9/2022 4:36:13 PM Unit 3.indd 79 0 IA IIA III B IV B VB VI B VII B III A IV A VA VI A VII A ← VIII → IB II B 2024-25 Classification of Elements and Periodicity in Properties 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 the elements. 79 9/9/2022 4:36:13 PM  80 chemistry a few atoms of them are obtained. Their digits which make up the atomic number and synthesis and characterisation, therefore, “ium” is added at the end. The IUPAC names require highly sophisticated costly equipment for elements with Z above 100 are shown in and laboratory. Such work is carried out with Table 3.5. competitive spirit only in some laboratories in the world. Scientists, before collecting the Table 3.4 Notation for IUPAC reliable data on the new element, at times Nomenclature of Elements get tempted to claim for its discovery. For example, both American and Soviet scientists Digit Name Abbreviation claimed credit for discovering element 104. 0 nil n The Americans named it Rutherfordium 1 un u whereas Soviets named it Kurchatovium. To 2 bi b avoid such problems, the IUPAC has made 3 tri t recommendation that until a new element’s 4 quad q discovery is proved, and its name is officially 5 pent p recognised, a systematic nomenclature be 6 hex h derived directly from the atomic number of 7 sept s the element using the numerical roots for 8 oct o 0 and numbers 1-9. These are shown in 9 enn e Table 3.4. The roots are put together in order of Table 3.5 Nomenclature of Elements with Atomic Number Above 100 Atomic Name according to IUPAC IUPAC Symbol Number IUPAC nomenclature Official Name Symbol 101 Unnilunium Unu Mendelevium Md 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 112 Ununbium Uub Copernicium Cn 113 Ununtrium Uut Nihonium Nh 114 Ununquadium Uuq Flerovium Fl 115 Ununpentium Uup Moscovium Mc 116 Ununhexium Uuh Livermorium Lv 117 Ununseptium Uus Tennessine Ts 118 Ununoctium Uuo Oganesson Og 2024-25 Unit 3.indd 80 9/9/2022 4:36:13 PM  Classification of Elements and Periodicity in Properties 81 Thus, the new element first gets a be readily seen that the number of elements temporary name, with symbol consisting in each period is twice the number of atomic of three letters. Later permanent name orbitals available in the energy level that is and 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 has two elements — hydrogen (ls1) and helium (or state of the country) in which the element (ls2) when the first shell (K) is completed. The was discovered, or pay tribute to a notable second period (n = 2) starts with lithium and the scientist. As of now, elements with atomic third electron enters the 2s orbital. The next numbers up to 118 have been discovered. element, beryllium has four electrons and has Official names of all elements have been the electronic configuration 1s22s2. Starting 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 (2s22p6). Thus there are What would be the IUPAC name and 8 elements in the second period. The third symbol for the element with atomic period (n = 3) begins at sodium, and the added number 120? electron enters a 3s orbital. Successive filling of 3s and 3p orbitals gives rise to the third Solution 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 added electrons fill up the 4s orbital. Now the symbol and the name respectively you may note that before the 4p orbital is filled, are Ubn and unbinilium. filling up of 3d orbitals becomes energetically favourable and we come across the so called 3d 3.5 ELECTRONIC CONFIGURATIONS transition series of elements. This starts from OF ELEMENTS AND THE PERIODIC scandium (Z = 21) which has the electronic TABLE configuration 3d1 4s2. The 3d orbitals are filled In the preceding unit we have learnt that an at zinc (Z=30) with electronic configuration electron in an atom is characterised by a set 3d104s2. The fourth period ends at krypton of four quantum numbers, and the principal with the filling up of the 4p orbitals. Altogether quantum number (n ) defines the main energy we have 18 elements in this fourth period. The level known as shell. We have also studied fifth period (n = 5) beginning with rubidium about the filling of electrons into different is similar to the fourth period and contains subshells, also referred to as orbitals (s, p, the 4d transition series starting at yttrium d, f ) in an atom. The distribution of electrons (Z = 39). This period ends at xenon with the into orbitals of an atom is called its electronic filling up of the 5p orbitals. The sixth period configuration. An element’s location in the (n = 6) contains 32 elements and successive Periodic Table reflects the quantum numbers electrons enter 6s, 4f, 5d and 6p orbitals, in of the last orbital filled. In this section we the order — filling up of the 4f orbitals begins will observe a direct connection between the with cerium (Z = 58) and ends at lutetium electronic configurations of the elements and (Z = 71) to give the 4f-inner transition series the long form of the Periodic Table. which is called the lanthanoid series. The seventh period (n = 7) is similar to the sixth (a) Electronic Configurations in Periods period with the successive filling up of the The period indicates the value of n for the 7s, 5f, 6d and 7p orbitals and includes most outermost or valence shell. In other words, of the man-made radioactive elements. This successive period in the Periodic Table is period will end at the element with atomic associated with the filling of the next higher number 118 which would belong to the noble principal energy level (n = 1, n = 2, etc.). It can gas family. Filling up of the 5f orbitals after 2024-25 Unit 3.indd 81 9/9/2022 4:36:13 PM  82 chemistry actinium (Z = 89) gives the 5f-inner transition a theoretical foundation for the periodic series known as the actinoid series. The 4f- classification. The elements in a vertical column and 5f-inner transition series of elements of the Periodic Table constitute a group or are placed separately in the Periodic Table family and exhibit similar chemical behaviour. to maintain its structure and to preserve the This similarity arises because these elements principle of classification by keeping elements have the same number and same distribution with similar properties in a single column. of electrons in their outermost orbitals. We can classify the elements into four blocks viz., Problem 3.2 s-block, p-block, d-block and f-block How would you justify the presence depending on the type of atomic orbitals that of 18 elements in the 5th period of the are being filled with electrons. This is illustrated Periodic Table? in Fig. 3.3. We notice two exceptions to this Solution categorisation. Strictly, helium belongs to the s-block but its positioning in the p-block When n = 5, l = 0, 1, 2, 3. The order in which the energy of the available along with other group 18 elements is orbitals 4d, 5s and 5p increases is 5s justified because it has a completely filled < 4d < 5p. The total number of orbitals valence shell (1s2) and as a result, exhibits available are 9. The maximum number properties characteristic of other noble gases. of electrons that can be accommodated The other exception is hydrogen. It has only is 18; and therefore 18 elements are one s-electron and hence can be placed in there in the 5th period. group 1 (alkali metals). It can also gain an electron to achieve a noble gas (b) Groupwise Electronic Configurations arrangement and hence it can behave Elements in the same vertical column or similar to a group 17 (halogen family) group have similar valence shell electronic elements. Because it is a special case, we configurations, the same number of electrons shall place hydrogen separately at the top of in the outer orbitals, and similar properties. the Periodic Table as shown in Fig. 3.2 and For example, the Group 1 elements (alkali Fig. 3.3. We will briefly discuss the salient metals) all have ns1 valence shell electronic features of the four types of elements marked in configuration as shown below. the Periodic Table. More about these elements Atomic number Symbol Electronic configuration 3 Li 1s22s1 (or) [He]2s1 11 Na 1s22s22p63s1 (or) [Ne]3s1 19 K 1s22s22p63s23p64s1 (or) [Ar]4s1 37 Rb 1s22s22p63s23p63d104s24p65s1 (or) [Kr]5s1 55 Cs 1s22s22p63s23p63d104s24p64d105s25p66s1 (or) [Xe]6s1 87 Fr [Rn]7s1 Thus it can be seen that the properties of will be discussed later. During the description an element have periodic dependence upon of their features certain terminology has been its atomic number and not on relative atomic used which has been classified in section 3.7. mass. 3.6 ELECTRONIC CONFIGURATIONS 3.6.1 The s-Block Elements AND TY P E S OF E L E M E NTS : The elements of Group 1 (alkali metals) and s-, p-, d-, f- BLOCKS Group 2 (alkaline earth metals) which have The aufbau (build up) principle and the ns1 and ns2 outermost electronic configuration electronic configuration of atoms provide belong to the s-Block Elements. They are all 2024-25 Unit 3.indd 82 9/9/2022 4:36:13 PM Unit 3.indd 83 Nh Mc Ts Og 2024-25 Classification of Elements and Periodicity in Properties 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 ( ). 83 11/2/2022 4:28:19 PM  84 chemistry reactive metals with low ionization enthalpies. valence (oxidation states), paramagnetism and They lose the outermost electron(s) readily to oftenly used as catalysts. However, Zn, Cd and form 1+ ion (in the case of alkali metals) or 2+ Hg which have the electronic configuration, ion (in the case of alkaline earth metals). The (n-1) d10ns2 do not show most of the properties metallic character and the reactivity increase of transition elements. In a way, transition as we go down the group. Because of high metals form a bridge between the chemically reactivity they are never found pure in nature. active metals of s-block elements and the The compounds of the s-block elements, with less active elements of Groups 13 and 14 and the exception of those of lithium and beryllium thus take their familiar name “Transition are predominantly ionic. Elements”. 3.6.2 The p-Block Elements 3.6.4 The f-Block Elements (Inner-Transition Elements) The p-Block Elements comprise those belonging to Group 13 to 18 and these The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, together with the s-Block Elements are Ce(Z = 58) – Lu(Z = 71) and Actinoids, called the Representative Elements or Main Th(Z = 90) – Lr (Z = 103) are characterised by Group Elements. The outermost electronic the outer electronic configuration (n-2)f1-14 configuration varies from ns2np1 to ns2np6 (n-1)d0–1ns2. The last electron added to each in each period. At the end of each period is element is filled in f- orbital. These two series a noble gas element with a closed valence of elements are hence called the Inner- shell ns2np6 configuration. All the orbitals Transition Elements (f-Block Elements). in the valence shell of the noble gases are They are all metals. Within each series, the completely filled by electrons and it is very properties of the elements are quite similar. difficult to alter this stable arrangement by The chemistry of the early actinoids is the addition or removal of electrons. The more complicated than the corresponding noble gases thus exhibit very low chemical lanthanoids, due to the large number of reactivity. Preceding the noble gas family oxidation states possible for these actinoid are two chemically important groups of non- elements. Actinoid elements are radioactive. Many of the actinoid elements have been made metals. They are the halogens (Group 17) and only in nanogram quantities or even less by the chalcogens (Group 16). These two groups nuclear reactions and their chemistry is not of elements have highly negative electron fully studied. The elements after uranium are gain enthalpies and readily add one or two called Transuranium Elements. electrons respectively to attain the stable noble gas configuration. The non-metallic Problem 3.3 character increases as we move from left to The elements Z = 117 and 120 have not yet right across a period and metallic character been discovered. In which family/group increases as we go down the group. would you place these elements and also give the electronic configuration in 3.6.3 The d-Block Elements (Transition each case. Elements) These are the elements of Group 3 to 12 in Solution the centre of the Periodic Table. These are We see from Fig. 3.2, that element characterised by the filling of inner d orbitals with Z = 117, would belong to the by electrons and are therefore referred to as halogen family (Group 17) and the d-Block Elements. These elements have electronic configuration would be [Rn] 5f146d107s27p5. The element with Z = 120, the general outer electronic configuration will be placed in Group 2 (alkaline earth (n-1)d1-10ns0-2 except for Pd where its electronic metals), and will have the electronic configuration is 4d105s0.. They are all metals. configuration [Uuo]8s2. They mostly form coloured ions, exhibit variable 2024-25 Unit 3.indd 84 11/2/2022 4:28:19 PM  Classification of Elements and Periodicity in Properties 85 3.6.5 Metals, Non-metals and Metalloids Solution In addition to displaying the classification of elements into s-, p-, d-, and f-blocks, Metallic character increases down a group and decreases along a period as Fig. 3.3 shows another broad classification we move from left to right. Hence the of elements based on their properties. The order of increasing metallic character elements can be divided into Metals and is: P < Si < Be < Mg < Na. Non-Metals. Metals comprise more than 78% of all known elements and appear on 3.7 PERIODIC TRENDS IN PROPERTIES the left side of the Periodic Table. Metals are OF ELEMENTS usually solids at room temperature [mercury There are many observable patterns in the is an exception; gallium and caesium also physical and chemical properties of elements have very low melting points (303K and as we descend in a group or move across a 302K, respectively)]. Metals usually have high period in the Periodic Table. For example, melting and boiling points. They are good within a period, chemical reactivity tends to conductors of heat and electricity. They are be high in Group 1 metals, lower in elements malleable (can be flattened into thin sheets by towards the middle of the table, and increases hammering) and ductile (can be drawn into to a maximum in the Group 17 non-metals. wires). In contrast, non-metals are located at Likewise within a group of representative the top right hand side of the Periodic Table. metals (say alkali metals) reactivity increases In fact, in a horizontal row, the property of on moving down the group, whereas within a elements change from metallic on the left to group of non-metals (say halogens), reactivity non-metallic on the right. Non-metals are decreases down the group. But why do the usually solids or gases at room temperature properties of elements follow these trends? with low melting and boiling points (boron And how can we explain periodicity? To and carbon are exceptions). They are poor answer these questions, we must look into the conductors of heat and electricity. Most non- theories of atomic structure and properties metallic solids are brittle and are neither of the atom. In this section we shall discuss malleable nor ductile. The elements become the periodic trends in certain physical and more metallic as we go down a group; the chemical properties and try to explain them non-metallic character increases as one goes in terms of number of electrons and energy from left to right across the Periodic Table. levels. The change from metallic to non-metallic 3.7.1 Trends in Physical Properties character is not abrupt as shown by the thick zig-zag line in Fig. 3.3. The elements (e.g., There are numerous physical properties of silicon, germanium, arsenic, antimony and elements such as melting and boiling points, tellurium) bordering this line and running heats of fusion and vaporization, energy diagonally across the Periodic Table show of atomization, etc. which show periodic variations. However, we shall discuss the properties that are characteristic of both periodic trends with respect to atomic and metals and non-metals. These elements are ionic radii, ionization enthalpy, electron gain called Semi-metals or Metalloids. enthalpy and electronegativity. Problem 3.4 (a) Atomic Radius Considering the atomic number and You can very well imagine that finding the position in the periodic table, arrange size of an atom is a lot more complicated than the following elements in the increasing measuring the radius of a ball. Do you know order of metallic character : Si, Be, Mg, Na, P. why? Firstly, because the size of an atom (~ 1.2 Å i.e., 1.2 × 10–10 m in radius) is very 2024-25 Unit 3.indd 85 9/9/2022 4:36:14 PM  86 chemistry small. Secondly, since the electron cloud The atomic radii of a few elements are listed surrounding the atom does not have a sharp in Table 3.6. Two trends are obvious. We can boundary, the determination of the atomic explain these trends in terms of nuclear charge size cannot be precise. In other words, there and energy level. The atomic size generally is no practical way by which the size of an decreases across a period as illustrated in individual atom can be measured. However, Fig. 3.4(a) for the elements of the second an estimate of the atomic size can be made by period. It is because within the period the knowing the distance between the atoms in outer electrons are in the same valence shell the combined state. One practical approach to and the effective nuclear charge increases estimate the size of an atom of a non-metallic as the atomic number increases resulting in element is to measure the distance between the increased attraction of electrons to the two atoms when they are bound together nucleus. Within a family or vertical column by a single bond in a covalent molecule and of the periodic table, the atomic radius from this value, the “Covalent Radius” of the increases regularly with atomic number as element can be calculated. For example, the illustrated in Fig. 3.4(b). For alkali metals bond distance in the chlorine molecule (Cl2) and halogens, as we descend the groups, is 198 pm and half this distance (99 pm), is the principal quantum number (n) increases taken as the atomic radius of chlorine. For and the valence electrons are farther from metals, we define the term “Metallic Radius” the nucleus. This happens because the inner which is taken as half the internuclear energy levels are filled with electrons, which distance separating the metal cores in the serve to shield the outer electrons from the metallic crystal. For example, the distance pull of the nucleus. Consequently the size of between two adjacent copper atoms in solid copper is 256 pm; hence the metallic radius the atom increases as reflected in the atomic of copper is assigned a value of 128 pm. For radii. simplicity, in this book, we use the term Note that the atomic radii of noble gases Atomic Radius to refer to both covalent or are not considered here. Being monoatomic, metallic radius depending on whether the their (non-bonded radii) values are very element is a non-metal or a metal. Atomic large. In fact radii of noble gases should be radii can be measured by X-ray or other compared not with the covalent radii but with spectroscopic methods. the van der 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 Atom Atomic Atom Atomic (Group I) Radius (Group 17) Radius Li 152 F 64 Na 186 Cl 99 K 231 Br 114 Rb 244 I 133 Cs 262 At 140 2024-25 Unit 3.indd 86 9/9/2022 4:36:14 PM  Classification of Elements and Periodicity in Properties 87 Fig. 3.4 (a) Variation of atomic radius with atomic Fig. 3.4 (b) Variation of atomic radius with number across the second period atomic number for alkali metals and halogens (b) Ionic Radius cation with the greater positive charge will The removal of an electron from an atom have a smaller radius because of the greater results in the formation of a cation, whereas attraction of the electrons to the nucleus. gain of an electron leads to an anion. The Anion with the greater negative charge will ionic radii can be estimated by measuring have the larger radius. In this case, the net the distances between cations and anions repulsion of the electrons will outweigh the in ionic crystals. In general, the ionic radii nuclear charge and the ion will expand in size. of elements exhibit the same trend as the atomic radii. A cation is smaller than its Problem 3.5 parent atom because it has fewer electrons Which of the following species will have while its nuclear charge remains the same. the largest and the smallest size? The size of an anion will be larger than that of Mg, Mg2+, Al, Al3+. the parent atom because the addition of one Solution or more electrons would result in increased Atomic radii decrease across a period. repulsion among the electrons and a decrease Cations are smaller than their parent in effective nuclear charge. For example, the atoms. Among isoelectronic species, ionic radius of fluoride ion (F–) is 136 pm the one with the larger positive nuclear whereas the atomic radius of fluorine is only charge will have a smaller radius. 64 pm. On the other hand, the atomic radius Hence the largest species is Mg; the of sodium is 186 pm compared to the ionic smallest one is Al3+. 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 A quantitative measure of the tendency of them isoelectronic species*. For example, an element to lose electron is given by its O2–, F–, Na+ and Mg2+ have the same number Ionization Enthalpy. It represents the of electrons (10). Their radii would be different energy required to remove an electron from an because of their different nuclear charges. The isolated gaseous atom (X) in its ground state. * Two or more species with same number of atoms, same number of valence electrons and same structure, regardless of the nature of elements involved. 2024-25 Unit 3.indd 87 9/9/2022 4:36:14 PM  88 chemistry In other words, the first ionization enthalpy for an 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 electron; it is the energy required to carry out the reaction shown in equation 3.2. X+(g) → X2+(g) + e– (3.2) Energy is always required to remove electrons from an atom and hence ionization Fig. 3.5 Variation of first ionization enthalpies (∆iH) with atomic number for elements enthalpies are always positive. The second with Z = 1 to 60 ionization enthalpy will be higher than the first ionization enthalpy because it is more can be correlated with their high reactivity. difficult to remove an electron from a positively In addition, you will notice two trends the charged ion than from a neutral atom. In the first ionization enthalpy generally increases same way the third ionization enthalpy will be as we go across a period and decreases higher than the second and so on. The term as we descend in a group. These trends “ionization enthalpy”, if not qualified, is taken are illustrated in Figs. 3.6(a) and 3.6(b) as the first ionization enthalpy. respectively for the elements of the second The first ionization enthalpies of elements period and the first group of the periodic having atomic numbers up to 60 are plotted table. You will appreciate that the ionization in Fig. 3.5. The periodicity of the graph is enthalpy and atomic radius are closely related quite striking. You will find maxima at the properties. To understand these trends, we noble gases which have closed electron shells have to consider two factors : (i) the attraction and very stable electron configurations. On of electrons towards the nucleus, and (ii) the the other hand, minima occur at the alkali repulsion of electrons from each other. The metals and their low ionization enthalpies effective nuclear charge experienced by a 3.6 (a) 3.6 (b) Fig. 3.6(a) F i r s t i o n i z a t i o n e n t h a l p i e s ( ∆ i H ) o f e l e m e n t s o f t h e s e c o n d p e r i o d a s a function of atomic number (Z) and Fig. 3.6(b) ∆iH of alkali metals as a function of Z. 2024-25 Unit 3.indd 88 9/9/2022 4:36:14 PM  Classification of Elements and Periodicity in Properties 89 valence electron in an atom will be less than the 2s electrons of beryllium. Therefore, it is the actual charge on the nucleus because of easier to remove the 2p-electron from boron “shielding” or “screening” of the valence compared to the removal of a 2s- electron from electron from the nucleus by the intervening beryllium. Thus, boron has a smaller first core electrons. For example, the 2s electron ionization enthalpy than beryllium. Another in lithium is shielded from the nucleus by “anomaly” is the smaller first ionization the inner core of 1s electrons. As a result, the enthalpy of oxygen compared to nitrogen. This valence electron experiences a net positive arises because in the nitrogen atom, three charge which is less than the actual charge 2p-electrons reside in different atomic orbitals of +3. In general, shielding is effective when (Hund’s rule) whereas in the oxygen atom, the orbitals in the inner shells are completely two of the four 2p-electrons must occupy the filled. This situation occurs in the case of same 2p-orbital resulting in an increased alkali metals which have single outermost electron-electron repulsion. Consequently, ns-electron preceded by a noble gas electronic it is easier to remove the fourth 2p-electron configuration. from oxygen than it is, to remove one of the When we move from lithium to fluorine three 2p-electrons from nitrogen. across the second period, successive electrons are added to orbitals in the same principal Problem 3.6 quantum level and the shielding of the nuclear The first ionization enthalpy (∆i H ) values charge by the inner core of electrons does of the third period elements, Na, Mg and not increase very much to compensate for Si are respectively 496, 737 and 786 kJ the increased attraction of the electron to the mol–1. Predict whether the first ∆i H value nucleus. Thus, across a period, increasing for Al will be more close to 575 or 760 kJ nuclear charge outweighs the shielding. mol–1 ? Justify your answer. Consequently, the outermost electrons are Solution held more and more tightly and the ionization It will be more close to 575 kJ mol–1. enthalpy increases across a period. As we go The value for Al should be lower than down a group, the outermost electron being that of Mg because of effective shielding increasingly farther from the nucleus, there is of 3p electrons from the nucleus by an increased shielding of the nuclear charge 3s-electrons. by the electrons in the inner levels. In this case, increase in shielding outweighs the (d) Electron Gain Enthalpy increasing nuclear charge and the removal of When an electron is added to a neutral the outermost electron requires less energy gaseous atom (X) to convert it into a negative down a group. ion, the enthalpy change accompanying the From Fig. 3.6(a), you will also notice that process is defined as the Electron Gain the first ionization enthalpy of boron (Z = 5) Enthalpy (∆ egH). Electron gain enthalpy is slightly less than that of beryllium (Z = 4) provides a measure of the ease with which even though the former has a greater nuclear an atom adds an electron to form anion as charge. When we consider the same principal represented by equation 3.3. quantum level, an s-electron is attracted to the nucleus more than a p-electron. In beryllium, X(g) + e – → X –(g) (3.3) the electron removed during the ionization is Depending on the element, the process an s-electron whereas the electron removed of adding an electron to the atom can be during ionization of boron is a p-electron. The either endothermic or exothermic. For many penetration of a 2s-electron to the nucleus is elements energy is released when an electron more than that of a 2p-electron; hence the 2p is added to the atom and the electron gain electron of boron is more shielded from the enthalpy is negative. For example, group nucleus by the inner core of electrons than 17 elements (the halogens) have very high 2024-25 Unit 3.indd 89 9/9/2022 4:36:15 PM  90 chemistry Table 3.7 Electron Gain Enthalpies* / (kJ mol–1) of Some Main Group Elements Group 1 ∆egH Group 16 ∆egH Group 17 ∆egH Group 0 ∆egH 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 Rb – 47 Te – 190 I – 295 Xe + 77 Cs – 46 Po – 174 At – 270 Rn + 68 negative electron gain enthalpies because they can attain stable noble gas electronic Problem 3.7 configurations by picking up an electron. Which of the following will have the most On the other hand, noble gases have large negative electron gain enthalpy and positive electron gain enthalpies because the which the least negative? electron has to enter the next higher principal P, S, Cl, F. quantum level leading to a very unstable Explain your answer. electronic configuration. It may be noted that Solution electron gain enthalpies have large negative values toward the upper right of the periodic Electron gain enthalpy generally table preceding the noble gases. becomes more negative across a period as we move from left to right. The variation in electron gain enthalpies of Within a group, electron gain enthalpy elements is less systematic than for ionization becomes less negative down a group. enthalpies. As a general rule, electron gain However, adding an electron to the enthalpy becomes more negative with increase 2p-orbital leads to greater repulsion in the atomic number across a period. The than adding an electron to the larger effective nuclear charge increases from left to 3p-orbital. Hence the element with right across a period and consequently it will most negative electron gain enthalpy is be easier to add an electron to a smaller atom chlorine; the one with the least negative since the added electron on an average would electron gain enthalpy is phosphorus. be closer to the positively charged nucleus. We should also expect electron gain enthalpy to (e) Electronegativity become less negative as we go down a group A qualitative measure of the ability of an atom because the size of the atom increases and the added electron would be farther from the in a chemical compound to attract shared nucleus. This is generally the case (Table electrons to itself is called electronegativity. 3.7). However, electron gain enthalpy of O or Unlike ionization enthalpy and electron gain F is less negative than that of the succeeding enthalpy, it is not a measureable quantity. element. This is because when an electron is However, a number of numerical scales of added to O or F, the added electron goes to electronegativity of elements viz., Pauling the smaller n = 2 quantum level and suffers scale, Mulliken-Jaffe scale, Allred-Rochow significant repulsion from the other electrons scale have been developed. The one which present in this level. For the n = 3 quantum is the most widely used is the Pauling scale. level (S or Cl), the added electron occupies Linus Pauling, an American scientist, in 1922 a larger region of space and the electron- assigned arbitrarily a value of 4.0 to fluorine, electron repulsion is much less. the element considered to have the greatest * 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. 2024-25 Unit 3.indd 90 9/9/2022 4:36:15 PM  Classification of Elements and Periodicity in Properties 91 ability to attract electrons. Approximate On the same account electronegativity values values for the electronegativity of a few decrease with the increase in atomic radii elements are given in Table 3.8(a) down a group. The trend is similar to that of ionization enthalpy. The electronegativity of any given element is not constant; it varies depending on the Knowing the relationship between element to which it is bound. Though it is electronegativity and atomic radius, can not a measurable quantity, it does provide a you now visualise the relationship between means of predicting the nature of force that electronegativity and non-metallic properties? holds a pair of atoms together – a relationship Non-metallic elements have strong tendency that you will explore later. Electronegativity generally increases across a period from left to right (say from lithium to fluorine) and decrease down a group (say from fluorine to astatine) in the periodic table. How can these trends be explained? Can the electronegativity be related to atomic radii, which tend to decrease across each period from left to right, but increase down each group ? The attraction between the outer (or valence) electrons and the nucleus increases as the atomic radius decreases in a period. The electronegativity also increases. 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 Atom Electronegativity Atom Electronegativity (Group I) Value (Group 17) Value 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 2024-25 Unit 3.indd 91 9/9/2022 4:36:15 PM  92 chemistry to gain electrons. Therefore, electronegativity is with outer electronic configuration 2s22p5, directly related to that non-metallic properties shares one electron with oxygen in the OF2 of elements. It can be further extended to say molecule. Being highest electronegative that the electronegativity is inversely related element, fluorine is given oxidation state to the metallic properties of elements. Thus, –1. Since there are two fluorine atoms in the increase in electronegativities across this molecule, oxygen with outer electronic a period is accompanied by an increase configuration 2s22p4 shares two electrons in non-metallic properties (or decrease in with fluorine atoms and thereby exhibits metallic properties) of elements. Similarly, the oxidation state +2. In Na2O, oxygen being decrease in electronegativity down a group is more electronegative accepts two electrons, accompanied by a decrease in non-metallic one from each of the two sodium atoms and, properties (or increase in metallic properties) thus, shows oxidation state –2. On the other of elements. hand sodium with electronic configuration All these periodic trends are summarised 3s1 loses one electron to oxygen and is given in Figure 3.7. oxidation state +1. Thus, the oxidation state of an element in a particular compound can 3.7.2 Periodic Trends in Chemical be defined as the charge acquired by its atom Properties on the basis of electronegative consideration Most of the trends in chemical properties of from other atoms in the molecule. elements, such as diagonal relationships, inert pair effect, effects of lanthanoid contraction Problem 3.8 etc. will be dealt with along the discussion Using the Periodic Table, predict the of each group in later units. In this section formulas of compounds which might we shall study the periodicity of the valence be formed by the following pairs of state shown by elements and the anomalous elements; (a) silicon and bromine properties of the second period elements (from (b) aluminium and sulphur. lithium to fluorine). Solution (a) Periodicity of Valence or Oxidation (a) Silicon is group 14 element with States a valence of 4; bromine belongs to The valence is the most characteristic property the halogen family with a valence of the elements and can be understood in of 1. Hence the formula of the compound formed would be SiBr4. terms of their electronic configurations. The valence of representative elements is usually (b) Aluminium belongs to group (though not necessarily) equal to the number 13 with a valence of 3; sulphur belongs to group 16 elements with of electrons in the outermost orbitals and/or a valence of 2. Hence, the formula equal to eight minus the number of outermost of the compound formed would be electrons as shown below. Al2S3. Nowadays the term oxidation state is frequently used for valence. Consider the Some periodic trends observed in the two oxygen containing compounds: OF2 and valence of elements (hydrides and oxides) Na2O. The order of electronegativity of the are shown in Table 3.9. Other such periodic three elements involved in these compounds trends which occur in the chemical behaviour is F > O > Na. Each of the atoms of fluorine, of the elements are discussed elsewhere in Group 1 2 13 14 15 16 17 18 Number of valence 1 2 3 4 5 6 7 8 electron alence 1 2 3 4 3,5 2,6 1,7 0,8 2024-25 Unit 3.indd 92 9/9/2022 4:36:15 PM  Classification of Elements and Periodicity in Properties 93 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 of LiH CaH2 B2H6 CH4 NH3 H2O HF hydride NaH AlH3 SiH4 PH3 H2S HCl KH GeH4 AsH3 H2Se HBr SnH4 H2Te HI Formula Li2O MgO B2O3 CO2 N2O3, N2O5 – of oxide Na2O CaO Al2O3 SiO2 P4O6, P4O10 SO3 Cl2 O7 K2O SrO Ga2O3 GeO2 As2O3, As2O5 SeO3 – BaO In2O3 SnO2 Sb2O3, Sb2O5 TeO3 – PbO2 Bi2O3 – – this book. There are many elements which the second element of the following group exhibit variable valence. This is particularly i.e., magnesium and aluminium, respectively. characteristic of transition elements and This sort of similarity is commonly referred actinoids, which we shall study later. to as diagonal relationship in the periodic properties. (b) Anomalous Properties of Second Period Elements What are the reasons for the different chemical behaviour of the first member of The first element of each of the groups 1 a group of elements in the s- and p-blocks (lithium) and 2 (beryllium) and groups 13-17 compared to that of the subsequent members (boron to fluorine) differs in many respects in the same group? The anomalous behaviour from the other members of their respective is attributed to their small size, large charge/ group. For example, lithium unlike other radius ratio and high electronegativity of the alkali metals, and beryllium unlike other elements. In addition, the first member of alkaline earth metals, form compounds with group has only four valence orbitals (2s and pronounced covalent character; the other 2p) available for bonding, whereas the second members of these groups predominantly member of the groups have nine valence form ionic compounds. In fact the behaviour orbitals (3s, 3p, 3d). As a consequence of of lithium and beryllium is more similar with this, the maximum covalency of the first member of each group is 4 (e.g., boron Property Element can only form  BF4  , whereas the other  members of the groups can expand their Metallic radius M/pm Li Be B valence shell to accommodate more than 152 111 88 four pairs of electrons e.g., aluminium Na Mg Al  AlF  3 6 forms). Furthermore, the first 186 160 143 member of p-block elements displays greater ability to form pπ – pπ multiple Ionic radius M+/pm Li Be bonds to itself (e.g., C = C, C ≡ C, 76 31 N = N, N ≡ Ν) and to other second period Na Mg elements (e.g., C = O, C = N, C ≡ N, 102 72 N = O) compared to subsequent members of the same group. 2024-25 Unit 3.indd 93 9/9/2022 4:36:15 PM  94 chemistry here it can be directly related to the metallic Problem 3.9 and non-metallic character of elements. Thus, Are the oxidation state and covalency of the metallic character of an element, which Al in [AlCl(H2O)5]2+ same ? is highest at the extremely left decreases and Solution the non-metallic character increases while No. The oxidation state of Al is +3 and moving from left to right across the period. the covalency is 6. The chemical reactivity of an element can be best shown by its reactions with oxygen and 3.7.3 Periodic Trends and Chemical halogens. Here, we shall consider the reaction Reactivity of the elements with oxygen only. Elements We have observed the periodic trends in on two extremes of a period easily combine certain fundamental properties such as with oxygen to form oxides. The normal oxide atomic and ionic radii, ionization enthalpy, formed by the element on extreme left is the electron gain enthalpy and valence. We most basic (e.g., Na2O), whereas that formed know by now that the periodicity is related to by the element on extreme right is the most electronic configuration. That is, all chemical acidic (e.g., Cl2O7). Oxides of elements in the and physical properties are a manifestation of centre are amphoteric (e.g., Al2O3, As2O3) or the electronic configuration of elements. We neutral (e.g., CO, NO, N2O). Amphoteric oxides shall now try to explore relationships between behave as acidic with bases and as basic with these fundamental properties of elements with acids, whereas neutral oxides have no acidic their chemical reactivity. or basic properties. The atomic and ionic radii, as we know, generally decrease in a period from left to right. Problem 3.10 As a consequence, the ionization enthalpies Show by a chemical reaction with water generally increase (with some exceptions as that Na2O is a basic oxide and Cl2O7 is an acidic oxide. outlined in section 3.7.1(a)) and electron gain enthalpies become more negative across a Solution period. In other words, the ionization enthalpy Na2O with water forms a strong base of the extreme left element in a period is the whereas Cl2O7 forms strong acid. least and the electron gain enthalpy of the Na2O + H2O → 2NaOH element on the extreme right is the highest Cl2O7 + H2O → 2HClO4 negative (note : noble gases having completely filled shells have rather positive electron Their basic or acidic nature can be gain enthalpy values). This results into high qualitatively tested with litmus paper. chemical reactivity at the two extremes and the lowest in the centre. Thus, the maximum Among transition metals (3d series), the chemical reactivity at the extreme left (among change in atomic radii is much smaller as alkali metals) is exhibited by the loss of an compared to those of representative elements electron leading to the formation of a cation across the period. The change in atomic radii and at the extreme right (among halogens) is still smaller among inner-transition metals shown by the gain of an electron forming (4f series). The ionization enthalpies are an anion. This property can be related with intermediate between those of s- and p-blocks. the reducing and oxidizing behaviour of the As a consequence, they are less electropositive elements which you will learn later. However, than group 1 and 2 metals.