PHYSICS-15-10- 11th (PQRS) SOLUTION

https://docs.google.com/document/d/1Sx-3HpGkd_pJqT0ErrRTkAKuZL5P_G-m/edit?usp=sharing&ouid=109474854956598892099&rtpof=true&sd=true Classification of Elements and Periodicity in Properties Modern periodic law and present form of the periodic table, s, p, d and f block elements, periodic trends in properties of elements, atomic and ionic radii, ionization enthalpy, electron gain enthalpy, valence, oxidation states and chemical reactivity. C H A P T E R GENESIS OF PERIODIC CLASSIFICATION Dobereiner grouped elements in triads e.g. Li, Na, K or Cl, Br, I etc. Newland found that similar elements are repeated at 8th place. It happened when elements were arranged in the increasing order of their atomic weight. It is applicable for lighter element only. Lother Mayer plotted a graph between atomic volume of elements against their atomic weight. He found that similar elements occupied similar positions on the curve. Mendeleev’s Periodic Law and Table: Mendeleev arranged all the elements in order of their increasing atomic weights. A table which has been formed with the help of classification of elements is called periodic table. The method of arranging similar elements in one group and separating them from dissimilar element is called classification of elements. Mendeleev prepared the table on the basis of his periodic law called Mendeleev’s periodic law. Mendeleev’s periodic law : The physical and chemical properties of elements are the periodic function of their atomic weights. Mendeleev’s periodic table consists of seven horizontal rows known as periods and nine vertical columns known as groups. Period : Out of the seven periods in all, the first three periods are known as short periods while the fourth, fifth and sixth periods are called long periods. THIS CHAPTER INCLUDES Genesis of Periodic Classification Nomenclature of superheavy elements Modern Periodic law Classification of Elements Periodic trends in properties Effective Nuclear Charge Atomic size Ionization enthalpy El ec tron affinity and Electron gain enthalpy Electronegativity Periods Period Total No. of Elements Starts with Elements Ends with Elements Remark 1. 2 Hydrogen (1) Helium (2) Very short period 2. 8 Lithium (3) Neon (10) Short period 3. 8 Sodium (11) Argon (18) Short period 4. 18 Potassium (19) Krypton (36) Long period 5. 18 Rubidium (37) Xenon (54) Long period 6. 32 Cesium (55) Radon (86) Very long period 7. 26 Francium (87) (Not named yet) (112) Incomplete period Groups There are nine groups in all including 8th group of transition elements and zero group of inert gases. All the group from I to VII (except zero and VIII) are divided into sub-groups A and B. The group number of an element represent its valency. The elements of same sub-group resemble one another more closely and generally differ to some extent from the elements of the other subgroups. Defects in Mendeleev’s Periodic Table Position of Hydrogen : Hydrogen resembles both the alkali metal (I group) and the halogen (VII). Hence its position in periodic table is undecided. Position of Isotopes :According to Mendeleev’s periodic law isotope of an element should occupy different position in the periodic table, but this is not so. Position of VIII group elements : Nine elements in the VIII group do not fit into the system. Position of Lanthanides and Actinides : Their position can not be justified according to the periodic law and cannot be arranged in the order of their increasing atomic weight. Dissimilar elements placed in the same group : Alkali metals (Li, Na, K, etc.) are placed with coinage metals (Cu, Ag, Au). Similar elements are placed apart : Chemically similar elements like Cu, Hg and Ag and Ti, Au and Pt have been placed in different groups. Anomalous pair of elements : Some elements of higher atomic weight have been placed before elements of lower atomic weight. For example, argon (At. wt. = 39.9) has been placed before potassium (At.wt = 39.1); cobalt (At. wt. = 58.94) is placed before nickel (At. wt. = 58.69) ; tellurium (127.5) has been placed before iodine (126.9). Modern periodic table : It is also known as long form or Bohr’s table as it is based on Bohr’s scheme of the arrangement of elements into four types according to their electronic configuration. Recent work has established that the fundamental property of an atom is atomic number and not the atomic weight. Therefore, atomic number is taken as the basis of the classification of elements. The modern periodic law may be stated as “The properties of elements are periodic function of their atomic number.” IUPAC nomenclature for the superheavy elements Atomic Number Name Symbol Atomic Number Name Symbol 110 un-un-nilium Uun 101 un-nil-unium Unu 111 un-un-unium Uuu 102 un-nil-bium Unb 112 un-un-bium Uub 103 un-nil-trium Unt 113 un-un-trium Uut 104 un-nil-quadium Unq 114 un-un-quadium Uuq 105 un-nil-pentium Unp 115 un-un-pentium Uup 106 un-nil-hexium Unh 116 un-un-hexium Uuh 107 un-nil-septium Uns 117 un-un-septium Uus 108 un-nil-octium Uno 118 un-un-octium Uuo 109 un-nil-ennium Une 119 un-un-ennium Uue 120 un-bi-nilium Ubn 130 un-tri-nilium Utn 140 un-quad-nilium Uqn 150 un-pent-nilium Upn CLASSIFICATION OF ELEMENTS On the basis of electronic configuration the elements can be classified into the following four types : s-block elements : These elements contain 1 or 2 electrons in s-subshell of outermost shell. Elements of 1 and 2 group belong to this class. These elements enter into chemical reaction by losing valency electrons so as to acquire noble gas configuration in the outermost orbit. ns1 (group 1) ns2 (group 2) (alkali metals) (alkaline earth metals) These elements generally form electrovalent compounds and basic oxides. p–block elements : These elements contain 1 to 6 electrons in the p–subshell of the outermost orbit (ns2 np1–6). The elements belonging to 13th to 18th group except He are p-block elements. In these last electron enters to the p-sub-shell. For example. 13 Boron (B) Z = 5 1s2 2s2 2p1 14 Carbon (C) Z = 6 1s2 2s2 2p2 15 Nitrogen (N) Z = 7 1s2 2s2 sp3 16 Oxygen (O) Z = 8 1s2 2s2 sp4 17 Fluorine (F) Z = 9 1s2 2s2 2p5 The main characteristics of these elements are : The non-metallic character increases along a period from 13 to 17. They form covalent compounds among themselves but electrovalent compounds with s-block elements. Their oxides are generally acidic, few are amphoteric also. For example Al2O3, Ga2O3. d–block elements : These are called transition elements or ‘d’ block elements. The elements of group 3 to 12 belong to this class. Their general configuration can be represented as : (n–1)d1–10 ns1–2 General characteristics of transition (d–block) elements: They are metals, hard, malleable, ductile and possess high tensile strength. They are good conductors of heat and electricity. These elements exhibit variable valency. They generally form coloured compounds. This is due to the presence of incomplete d–subshell. These metals, their alloys and compounds possess marked catalytic activity. They are generally paramagnetic, i.e., attract magnetic lines of force. f–block elements : They are inner transition or f-block elements. These elements are arranged in the two row at the bottom of the periodic table. In the f irst row 14 elements f rom atomic number 58 to 71, known as Lanthanides or rare earth elements. The second row of elements from atomic number 90 to 103, known as actinides. Their general electronic configuration can be represented as (n – 2) f 1–14 (n–1)d 0–1 ns2 They show most of the properties similar to each other since outermost and penultimate orbits are similar. Their properties are similar to ‘d’ block elements. PERIODIC TRENDS IN PROPERTIES Atomic Size (or atomic radius) Atomic radius is the size of the atom of an element. Atomic radius is defined as “the distance from the centre of the nucleus upto the centre of outermost electron.” It is measured in Angstrom unit (Å). It is not possible to measure exact atomic radius as an atom is unstable and it cannot be isolated to get its radius. Moreover, the exact position of the outermost electron is uncertain. The values for radii are obtained from x-ray measurements. Following points are to be noted in this reference : The size of an atom or ion decreases in a horizontal period as we move from left to right. The atomic radius increases in a group with the rise in atomic number. A positive ion (cation) is smaller than the corresponding atom : A positive ion or cation is formed by the loss of one or more electrons from an atom and the number of protons remains the same in the nucleus. Thus the ratio of the positive charge in the nucleus to the number of electrons i.e., effective nuclear charge increases. Hence the force of attraction of nucleus to the outer electrons increases thus decreasing the size of cation. In case of alkali metals, the removal of an electron removes the entire outermost shell. A negative ion (anion) is bigger than the corresponding atom : In the formation of negative ion (anion) one or more electrons(s) are added to the atom. Thus results in the expansion of the size of the nuclear charge, which in turn decreases the force of attraction and increases the size of an anion or the pull exercised by the nucleus on the electron become less i.e., they move a little farther resulting in an increase in the ionic size. Ionization Enthalpy It is the amount of energy required to remove most loosely held electron from the ground state of an Avogadro number of the isolated atoms, ions or molecules in the gaseous state. The ion formed by loss of first electron may lose further electrons and thus we may have successive ionization energies for removal of 2nd, 3rd and 4th electrons in the gaseous state. Ionization is always an endothermic process and ionization energies are therefore always assigned positive values. Factors influencing Ionization energy Successive Ionization - Generally ionization energy increases for successive ionizations. Atomic size - Ionization energy decreases as the size of atom increases. Value of Z - Higher the value of Z, higher is the I.E. Distance of electron from the nucleus - Smaller the distance of the electron from the nucleus larger is the ionization energy Sheilding effect - Higher is the sheilding of the electron to be removed lower is the I.E. Sheilding effect of the electrons of different orbitals follows the order s > p > d > f. Penetration effect - Higher the penetration power of the electron to be removed higher is the I.E. The penetration power of electrons of various orbitals follow the order s > p > d > f. Nature of shell - Ionization energy increases if the electron to be ionized from the species belongs to a half filled shell or a completely filled shell. The relative stability of the these configurations follows the order d5 < p3 < d10 M – Cl > M – Br > M – I NaF > NaCl > NaBr > NaI The order of melting point of chlorides of alkali metals is as follows : LiCl < CsCl < RbCl < KCl < NaCl The melting point of LiCl is lowest because it is with highest covalent character. The solubility of alkali metal carbonates in water at 298 K increases down the group from Lithium to Cesium. The basic character of oxides and hydroxides of group 1 and group 2 increases down the group because metallic character increases down the group e.g., LiOH is least basic whereas CsOH is most basic. Be(OH)2 is amphoteric, Mg(OH)2 is a weak base, Ca(OH)2 and Sr(OH)2 are moderately strong bases, Ba(OH)2 is strong base. The solubility of hydroxides of Group 1 and Group 2 in water increases down the group. The solubility of sulphates, carbonates and phosphates decreases down the Group 2 because lattice energy dominates over hydration energy in Group 2, for example MgSO4 is soluble in water wherease BaSO4 is insoluble in water. Li2CO3 is thermally unstable whereas other alkali metal carbonates are thermally stable. Thermal stability of carbonates of Group 2 increases down the group. All are thermally unstable. Properties of Li almost similar to that of Mg, Be are almost similar to that of Al and B are almost similar to that of C due to diagonal relationship. ❑ ❑ ❑