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The Periodic Table (very) Brief History • 1869 Mendeleev* & Meyer published similar tables *First to be recognized at international convention – Elements were ordered based on atomic mass • Henry Moseley developed the atomic # concept – Proved more accurate than Mendeleev’s atomic mass method The Periodic Law • When elements are arranged by increasing atomic # their chemical & physical properties show a periodic pattern. Review of Element Symbol • Atomic # – #protons • Element Symbol • Atomic Mass 54.938049 54.938049 – Weighted average of all isotopes’ mass # – Listed in AMUs – Equal to MM (g/mol) • Element Name Periodic Table Layout • Groups or Families – Columns, Vertical – 18 groups • Periods Groups to know – – – – Group 1 - Alkali Metals Group 2 - Alkaline Earth Metals Group 17 Halogens Group 18 - Noble Gases – Rows, Horizontal – 7 periods • Kinds of Elements – Metals, Nonmetals, Semi-metals – Varying Properties Electron Configurations Using the Periodic Table s-block elements: 1 Group 1: Alkali H- 1s1 Li - 1s22s1 H 2 IIA 3 4 Li Be 1.00797 2 3 Na- 1s22s22p63s1 K - 1s22s22p63s23p64s1 1 IA 1 4 5 6 7 6.939 11 9.0122 12 Na Mg 22.9898 24.305 19 20 K 39.102 37 Rb Ca 40.08 38 Sr 85.47 55 87.62 56 Cs Ba Fr Ra 132.905 137.34 87 88 [223] [226] s-block elements: Group 2: Alkaline Earth Be - 1s22s2 Mg- 1s22s22p63s2 Ca - 1s22s22p63s23p64s2 Electron Configurations (con’t) p-block elements: Group 13: B -1s22s22p1 Group 14: Si -1s22s22p63s23p2 Group 15: As-1s22s22p63s23p64s23d104p3 Group 16: S 1s22s22p63s23p4 Group 17: Br1s22s22p63s23p64s23d104p5 13 IIIA 14 IVA 15 VA 16 VIA 17 VIIA 18 VIIIA 2 He 5 6 7 8 9 4.0026 10 B C N O F Ne Al Si P S Cl Ar Ge As Se Br Kr In Sn Sb Te I Xe Tl Pb Bi Po At Rn 10.811 12.0112 14.0067 15.9994 18.9984 20.179 13 14 15 16 17 18 26.9815 28.086 30.9738 32.064 35.453 39.948 31 32 33 34 35 36 Ga 7 65.37 49 72.59 74.9216 78.96 79.909 83.80 50 51 52 53 54 40 114.82 118.69 121.75 127.60 126.904 131.30 81 82 83 84 85 86 59 204.37 207.19 208.980 [210] [210] [222] Noble Gas Configuration Choose Noble Gas Prior to Element FOCUS ON THE VALENCE SHELL Put Noble gas in brackets, ex. [Ne] This represents the “inner shell” of the element Add the outer level electrons by row # (2-7) s & p **d level starts on 4th row** 4f & 5f start on 6th & 7th row Examples: Ni: [Ar] 4s23d8 **remember that 3d fills after 4s and before 4p Sb: [Kr] 5s24d105p3 Trends in the Periodic Table Definition: Predictable changes in properties of the elements as you move through the table. (Realize there are exceptions) Decrease Increase Increase Increas e Atomic Radii 1 1 H 2 Li Be 3 Na Mg 11 12 4 K 5 Rb 6 Cs 7 Fr 3 19 37 55 87 21 Sc 38 39 56 Ba 88 Ra He 5 B - Ions are larger Ca Sr 2 Ionic Size +Ions are Smaller 4 20 Ionization Energy Electron Affinity Y 57 22 Ti 40 Zr 72 La Hf 89 104 Ac Rf 23 24 V Cr 41 42 25 43 Nb Mo Tc 73 44 27 Co 45 Ru Rh Ta W Re Os 75 76 Ir 105 106 107 108 109 Db 74 26 Mn Fe Sg 77 13 28 29 Ni Cu 46 47 Pd Ag 78 Pt 79 30 Zn 48 Cd 80 Au Hg 6 7 N O 16 Al Si 14 15 31 32 33 P Ga Ge As 49 In 81 Tl 8 C 50 Sn 82 Pb 51 Sb 83 Bi S 34 9 F 17 Cl 35 Se Br 52 53 Te 84 Po Bh Hs Mt Atomic Radii - distance from the nucleus to outermost electron. Ionization Energy - energy required to remove an electron (kJ/mol) Electron Affinity – energy change when neutral atom gains electron I 85 At 10 Ne 18 Ar 36 Kr 54 Xe 86 Rn Ionization Energy • In general, ionization energies of the main-group elements increase across each period. • This increase is caused by increasing nuclear charge. • A higher charge more strongly attracts electrons in the same energy level. • Among the main-group elements, ionization energies generally decrease down the groups. • Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus.(electron shielding) • The electrons are removed more easily. Additional Trend Information Electron Affinity The energy associated with an atom gaining or losing an electron. (kJ/mol) • + Energy means it requires energy …not favorable • - Energy means it gives up energy…favorable Electronegativity The ability to attract an electron during bonding • Increases up a group and across a period Electron Affinity • The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. • Electron affinity generally increases across periods. • Increasing nuclear charge along the same sublevel attracts electrons more strongly • Electron affinity generally decreases down groups. • The larger an atom’s electron cloud is, the farther away its outer electrons are from its nucleus. Ionic Radii • A positive ion is known as a cation. • The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius. • The electron cloud becomes smaller. • The remaining electrons are drawn closer to the nucleus by its unbalanced positive charge. • A negative ion is known as an anion. • The formation of an anion by the addition of one or more electrons always leads to an increase in atomic radius. Ionic Radii, continued • Cationic and anionic radii decrease across a period. • The electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level. • The outer electrons in both cations and anions are in higher energy levels as one reads down a group. • There is a gradual increase of ionic radii down a group. Electronegativity • Valence electrons hold atoms together in chemical compounds. • In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another. • Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. • Electronegativities tend to increase across periods, and decrease or remain about the same down a group. The Periodic Table (very) Brief History • 1800’s - Dobereiner introduced “Triads” – 3 elements with similar properties • 1865 - Newlands introduced “Law of Octaves” – At this time 62 known elements – 1st behaved like 8th, 2nd like 9th… • 1869 Mendeleev* & Meyer published similar tables *First to be recognized at international convention – Elements were ordered based on atomic mass The Periodic Law • When elements are arranged by increasing atomic # their chemical & physical properties show a periodic pattern. 1913 - Henry Moseley developed the atomic # concept. – Proved more accurate than Mendeleev’s atomic mass method Review of Element Symbol • Atomic # – #protons • Element Symbol • Atomic Mass 54.938049 – Weighted average of all isotopes’ mass # – Listed in AMUs – Equal to MM (g/mol) • Element Name Periodic Table Layout • Groups or Families – Columns, Vertical – 18 groups • Periods Groups to know – – – – Group 1 - Alkali Metals Group 2 - Alkaline Earth Metals Group 17 Halogens Group 18 - Noble Gases – Rows, Horizontal – 7 periods • Kinds of Elements – Metals, Nonmetals, Semi-metals – Varying Properties Electron Configurations Using the Periodic Table s-block elements: Group 1: Alkali H - 1s1 Li - 1s22s1 Na- 1s22s22p63s1 K- 1 1 IA 1 H 2 IIA 3 4 1.00797 2 3 4 5 1s22s22p63s23p64s1 6 7 Li 6.939 11 Na Be 9.0122 12 Mg s-block elements: Group 2: Alkaline Earth Be - 1s22s2 22.9898 24.305 19 20 Mg- 1s22s22p63s2 39.102 37 40.08 38 Ca - 1s22s22p63s23p64s2 85.47 55 87.62 56 K Rb Cs Ca Sr Ba 132.905 137.34 87 88 Fr [223] Ra [226] Electron Configurations (con’t) p-block elements: Group 13: B -1s22s22p1 13 IIIA 5 14 IVA 6 15 VA 7 16 17 VIA VIIA 8 9 18 VIIIA 2 He 4.0026 10 B C N O F Ne Al Si P S Cl Ar Ge As Se Br Kr Sn Sb Te I Xe Pb Bi Po At Rn Group 14: Si -1s22s22p63s23p2 10.811 12.0112 14.0067 15.9994 18.9984 20.179 13 14 15 16 17 18 Group 15: As-1s22s22p63s23p64s23d104p3 26.9815 28.086 30.9738 32.064 35.453 39.948 31 32 33 34 35 36 Ga 7 65.37 49 Group 16: S 1s22s22p63s23p4 In 40 114.82 118.69 121.75 127.60 126.904 131.30 81 82 83 84 85 86 Tl Group 17: Br- 72.59 74.9216 78.96 79.909 83.80 50 51 52 53 54 59 204.37 207.19 208.980 [210] 1s22s22p63s23p64s23d104p5 [210] [222] Noble Gas Configuration Choose Noble Gas Prior to Element FOCUS ON THE VALENCE SHELL Put Noble gas in brackets, ex. [Ne] This represents the “inner shell” of the element Add the outer level electrons by row # (2-7) s & p **d level starts on 4th row** 4f & 5f start on 6th & 7th row Examples: Ni: [Ar] 4s23d8 **remember that 3d fills after 4s and before 4p Sb: [Kr] 5s24d105p3 Trends in the Periodic Table Definition: Predictable changes in properties of the elements as you move through the table. Decrease Increase Increas e Atomic Radii 1 Ionization Energy 2 1 H 2 Li Be 4 +Ions are Smaller 3 11 12 Na Mg - Ions are larger 4 K 5 Rb Sr 6 Cs Ba La Hf Ta W Re Os Ir 7 Fr Ra Ac Rf 109 3 19 Ionic Size 20 21 37 38 39 Y Zr Nb Mo Tc 55 56 87 88 He 5 B 13 6 7 8 C N O Al Si 14 15 16 32 33 34 50 51 52 30 31 44 45 46 47 48 49 78 79 80 81 35 36 Cr Mn Fe Co Ni Cu Zn Ga Ge As Se 24 25 Br Kr 40 41 42 43 53 57 72 73 74 75 76 89 104 106 107 108 105 29 18 Ar V Db 28 17 Cl 23 Ti 27 S 10 Ne 22 Ca Sc 26 P 9 F Ru Rh Pd Ag Cd In 77 Pt Au Hg Tl Sn Sb Te 82 Pb 83 Bi 84 Po Sg Bh Hs Mt Atomic Radii - distance from the nucleus to outermost electron. Ionization Energy - energy required to remove an electron (kJ/mol) 54 I Xe 85 86 At Rn Additional Trend Information Electron Affinity The energy associated with an atom gaining or losing an electron. (kJ/mol) • + Energy means it requires energy …not favorable • - Energy means it gives up energy…favorable Electronegativity The ability to attract an electron during bonding • Increases up a group and across a period