RI 2022 Periodic Table I

' Raffles Institution Year 5 H2 Chemistry 2022 Lecture Notes 12 - The Periodic Table I Content • Periodicity of atomic and physical properties of the elements : variation with proton number across the third period (sodium to chlorine) of: (i) electronic configuration (ii) atomic radius and ionic radius (iii) ionisation energy (iv) electronegativity (v) melting point (vi) electrical conductivity • Periodicity of chemical properties of the elements in the third period: (i) variation in oxidation number and bonding of the oxides (sodium to sulfur only) and of the chlorides (sodium to phosphorus only) (ii) reactions of these oxides and chlorides with water (iii) acid/base behaviour of these oxides and the corresponding hydroxides Learning Outcomes Trends and variations in atom ic and physical properties For elements in the third period (sodium to chlorine), candidates should be able to: (a) recognise variation in the electronic configurations across a Period (b) describe and explain qualitatively the trends and variations in atomic radius, ionic radius, first ionisation energy and electronegativity: (i) across a Period in terms of shielding and nuclear charge· (c) interpret the variation in melting point and in electrical conductivity across a Period in terms of structure and bonding in the elements (metallic, giant molecular , or simple molecular) Trends and variations in chem ical ,1,r~!~-\::::f;·.: For elements in the third period (sc:'•:._;:·, k. ,:'1:0dne) candidates should be able to: (d) (i) state and explain the varic:t,u'.l ;;, me highest oxidation number of the elements in oxides (for Na20 ; MgO ; Al2Q3; Si0 2; 1-'.,0,,,; :,U:,) antl chlorides (for NaCl ; MgCh ; AlCfa; SiCl4; PCls) (ii) state and explain the varia.tic,n in bonding in oxides and chlorides in terms of electronegativity (with the exception of A/C'3) (iii) describe the reactions of the oxides with water (for Na20; MgO; Af2Q3; Si02; P4010; S03) (iv) describe and explain the acid/base behaviour of oxides (for Na20 ; MgO; Al203; Si02; P4010; S03) and hydroxides (for NaOH ; Mg(OH)2; A/(OH)3), including, where relevant, amphoteric behaviour in reaction with sodium hydroxide (only) and acids (v) describe and explain the reactions of the chlorides with water (for NaCl ; MgC!i; AlCfa; SiC/4; PCls) (vi) suggest the types of structure and bonding present in the oxides and chlorides from observations of their chemical and physical properties In addition, candidates should be able to: (e) predict the characteristic properties of an element in a given Group by using knowledge of chemical periodicity (f) deduce the nature , possible position in the Periodic Table , and identity of unknown elements from given information of physical and chemical properti es Lecture Outline 1 The Periodic Table 2 Trends and variations in atomic and physical References 1 Chemistry for advanced level (by Peter Cann and Peter Hughes) properties 2 Chemistry in Context (by Hill and Holman) Chemistry: The Molecular Nature of Matter and Change (by Martin S. Silberberg) 3 Trends and variations in chemical properties 3 4 Website : http://www .chemguide .co.uk -1-

1 THE PERIODIC TABLE The Periodic Table is the arrangement of the elements in increasing order of proton number such that elements with similar physical and chemical properties are grouped together. • Horizontal rows in the Periodic Table are called periods. - Elements in the same period have the same number of electronic shells. • Vertical columns are called groups. - Elements in the same group have the same outer electron configuration. Periodicity refers to the recurrence of similar properties at regular intervals when elements are arranged in increasing atomic number. Period 1 Period2 Period 3 Period4 Period 5 Period& Period 7 ns2 np1 - - - - - - - - - ♦ ns2 np5 3 4 5 6 7 8 10 11 12 13 14 15 16 17 18 &-block p-bloclt f-block lanthanoids I actinoids ..__ ____________________ ___, The elements are further grouped into four blocks. (i) s-block (electronic configuration of valence shell: ns1, ns2) • consists of metallic elements from Group 1 (alkali metals} and Group 2 (alkaline earth metals) and the non-metal, hydrogen • the metals of Group 1 and Group 2 have lower densities, melting points and boiling points than other metais • the metals are reactive , i.e. high in !he electrochemical series • the metals form stable , non--volatiie ionic compounds with non-metals (ii) p-block (electronic configuration of valence shell : ns2np1 to ns2np6) • electronic configuration of helium is 1 s2 • consists of elements from Group 13 to Group 18 • generally have low densities, melting points and boiling points (with exceptions, e.g. diamond) • generally form ionic compounds with metals and molecular compounds (which may be macromolecules or simple molecules) with other non-metals • noble gases are generally chemically inert (iii) d-block ( electronic configuration, e.g. [Ar)3d14s2 to [Ar]3d104s2) • includes transition metals • higher densities, melting points and boiling points than s-block metals • less reactive than s-block metals • marked similarities across a period, e.g. form coloured complexes (iv) f-block • lanthanoids and actinoids • starts from the sixth period -2- ' •

2 TRENDS AND VARIATIONS IN ATOMIC AND PHYSICAL PROPERTIES OF ELEMENTS 2.1 Variations in atomic properties Graph of atomic and Ionic raill against atomic number of Period 3 elements (N ■ to Cl) 0.25 ,------- - ------~ pl- ~ 0.20 --------,--.......,~---+ -• .:! ,, anionic radius f 0.15 t------'""-<::::;---~-----t 0 'E .!! :; 0.10 J.--------i-----l=-w atomic radius 0 e 0 ~ 0.05 1---AA11p.c..._::,...=--t-------t cationic radius s1•• 0.00 1---~-~-~-~-~--+ 11 12 Across a period, • the number of shells remain the same. • number of protons increases and hence nuclear charge increases. • number of electrons also increases but these electrons are added to the same outermost shell, and hence shielding effect remains approximately constant. • effective nuclear charge increases. • electrostatic attraction between the nucleus and the valence electrons increases, resulting in a decrease in the size of the electron cloud. Hence atomic radii decrease across a period. 13 14 15 16 17 Atomic number (i) The radius of a cation is always smaller than that of the parent atom. • Both the cation and its parent atom have the same number of protons and hence have the same nuclear charge. • However, the cation has one less electronic shell than its parent atom • Electrostatic attraction between the nucleus and the valence electrons increases, resulting in a decrease in size of the electron cloud. (ii) The radius of an anion is always greater than that of the parent atom. Both the anion and its parent atom have the same number of protons and hence have the same nuclear charge. • Hcwever, lhe anion has more electrons than its parent ·Ni!i•1 r::are electrons, electron-electron repulsion increases i_-,.•,drL,static attraction between the nucleus and the ·1.1i,_, r:c.:l electrons decreases, resulting in an increase in the ~i1.e of tile electron cloud. (iii) Ionic radii of isoelectronic ions decrease across a period. For example, from Na• to Si4• (or from P3- to Ct), • number of protons and hence nuclear charge increases while • their valence electrons experience the same shielding effect • effective nuclear charge increases • electrostatic attraction between the nucleus and the valence electrons increases, resulting in the decrease in the size of the electron cloud. Note: This reasoning can be extended to include other ions which are isoelectronic to Na ♦• i.e. to explain the decreasing ionic radii from C4- to Si4• . (similarly for p:,... to Ca2• ) -3-

Na Mg Al Si p s Cl (i) The first ionisation energies of the elements generally increase across a period. Ac