Periodic Table

1. Periodic Table

2. Russian chemist Organized elements by properties Arranged elements by atomic mass Predicted existence of several unknown elements Element 101 - Mendeleevium (Md) Dmitri Mendeleev Dmitri Mendeleev

3. Mendeleev’s Periodic Table

4. Elements Properties are Predicted

5. Modern Periodic Table • Based on Mendeleev’s system • Arrange elements by atomic number • Elements organized by properties • Columns are groups or families(1 to 18) • Rows are periods (1 to 7) showing the variation of chemical and physical properties.

6. Groups of Elements 1A 8A 1A 5A Nitrogen group Alkali metals H 1 2A 6A He 2 Alkali earth metals Oxygen group 1 1 2A 7A 3A 4A 5A 6A 7A Halogens Transition metals 3A 8A Noble gases Boron group Li 3 Be 4 B 5 C 6 N 7 O 8 F 9 Ne 10 2 2 4A Carbon group Hydrogen Inner transition metals Na 11 Mg 12 Al 13 Si 14 P 15 S 16 Cl 17 Ar 18 3 3 8B 3B 4B 5B 6B 7B 1B 2B K 19 Ca 20 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Ga 31 Ge 32 As 33 Se 34 Br 35 Kr 36 4 4 Rb 37 Sr 38 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48 In 49 Sn 50 Sb 51 Te 52 I 53 Xe 54 5 5 Cs 55 Ba 56 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Tl 81 Pb 82 Bi 83 Po 84 At 85 Rn 86 * * 6 6 Fr 87 Ra 88 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 W W 7 7 La 57 Ce 58 Pr 59 Nd 60 Pm 61 Sm 62 Eu 63 Gd 64 Tb 65 Dy 66 Ho 67 Er 68 Tm 69 Yb 70 Lu 71 * Ac 89 Th 90 Pa 91 U 92 Np 93 Pu 94 Am 95 Cm 96 Bk 97 Cf 98 Es 99 Fm 100 Md 101 No 102 Lr 103 W

7. Metals and Nonmetals H 1 He 2 1 Li 3 Be 4 B 5 C 6 N 7 O 8 F 9 Ne 10 Nonmetals 2 Na 11 Mg 12 Al 13 Si 14 P 15 S 16 Cl 17 Ar 18 3 K 19 Ca 20 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Ga 31 Ge 32 As 33 Se 34 Br 35 Kr 36 4 METALS Rb 37 Sr 38 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48 In 49 Sn 50 Sb 51 Te 52 I 53 Xe 54 5 Cs 55 Ba 56 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Tl 81 Pb 82 Bi 83 Po 84 At 85 Rn 86 * 6 Fr 87 Ra 88 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 W 7 Metalloids La 57 Ce 58 Pr 59 Nd 60 Pm 61 Sm 62 Eu 63 Gd 64 Tb 65 Dy 66 Ho 67 Er 68 Tm 69 Yb 70 Lu 71 Ac 89 Th 90 Pa 91 U 92 Np 93 Pu 94 Am 95 Cm 96 Bk 97 Cf 98 Es 99 Fm 100 Md 101 No 102 Lr 103

8. Properties of Metals, Nonmetals, and Metalloids METALS malleable, lustrous, ductile, good conductors of heat and electricity NONMETALS gases or brittle solids at room temperature, poor conductors of heat and electricity (insulators) METALLOIDS (Semi-metals) dull, brittle, semi-conductors (used in computer chips) exhibit properties of both metals and nonmetals

9. Electrons filling the orbitals shown in the Periodic Table 1 8 Groups 2 1s 1 3 4 5 6 7 1s 2s 2 2p 3s 3p 3 4p 3d Periods 4s 4 4d 5p 5s 5 La 5d 6p 6 6s Ac 6d 7 7s 4f Lanthanide series 5f Actinide series

10. Size of Atoms - Trends Periodic Trends in Atomic Radii

11. Relative Size of Atoms

12. Shielding Effect Valence Kernel electrons block the attractive force of the nucleus from the valence electrons + - - nucleus - Electrons - Electron Shield “kernel” electrons

13. Atomic Radius vs. Atomic Number

14. Atomic Radii trend explained • As you go across the period the number of shielding electrons are the same. • The nuclear charge is increasing (adding protons as you go across). • The electrons added are in the same valence shell – same distance from the nucleus. • More + nuclear charge gets out to the valence electrons, pulling the valence electrons in closer (stronger attraction). Ne Li Li As we go down a group each atom has another energy level, so the atoms get bigger. There are more levels in the kernel and therefore greater shielding of valence electrons (weaker attraction). Cs

15. Why do elements react? • Atoms with filled valence shells are stable – low in energy. Atoms attain a full valence shell by losing, gaining or sharing valence electrons. The result is a particle which is isoelectronic with a noble gas (has the same electron configuration as a noble gas). • Na 1+ =1s2 2s2 2p6 (isoelectronic Ne). • F1- = 1s2 2s2 2p6

16. Atomic and Ionic Radii vs Atomic Number

17. Radii period trends explained • Metals are “born losers”, the atoms lose their valence electrons to form cations (+ ions). The kernel is smaller. The remaining electrons are more strongly attracted to the nucleus. There are more protons than electrons so the remaining electrons are pulled in closer (+ ion is smaller). • Nonmetals will gain valence electrons to fill the valence shell to form anions (- ions). There are more electrons to share the nuclear charge so there is a weaker attraction between the electrons and the nucleus. The weaker attraction leads to the valence electrons being further away (- ion is larger) N 3 - O 2 - F 1 - Li1+ B3+ C4+ Be2+

18. Radii group trends explained Li+1 • As you go down the group you are adding energy levels so the cations and anions get bigger. Na+1 K+1 Rb+1 Cs+1

19. First Ionization Energy Plot

20. First ionization energy trend explained • 1st ionization energy is the energy required to remove one electron from the gaseous atom of an element. • 1st ionization energy is increased by strong attraction to valence electrons and a stable electron configuration. • As you go across the period the attraction to valence electrons increases so ionization energy increases. • Peaks in energy occur at Be (full s sublevel), N (½ full p sublevel) and Ne (full s & p sublevels –full valence shell).

21. Electronegativity values • Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element. • Big electronegativity values means the atom pulls the valence electrons toward the nucleus, strong attraction to valence electrons. • Atoms with small electronegativity values have weak attraction to valence electrons and these electrons are drawn away from this atom.

22. Electronegativity tend explained • Group Trend: The further down a group the farther the electron is away and the more shielding electrons an atom has – the lower the electronegativity value. Period Trend: Metals on the left lose electrons easily (weak attraction for valence electrons). Metals have low electronegativity values. • Nonmetals on the right need more electrons to complete the valence shell and have strong attraction for valence electrons. Nonmetals have high electronegativity values. • Electronegativity values increase as you go across the period

23. Electron Affinity values • Electron Affinity is the energy change associated with adding an electron to a gaseous atom. • If the atom becomes more stable (electron configuration like the noble gases) there is a loss of energy and the energy change is shown a negative value. • If the atom becomes less stable the energy level goes up and the energy change is shown as a positive value.

24. Electron Affinity tend explained • Across a period the electron affinity is low for metals and high for nonmetals (electron affinity increases from left to right). In Group 7A valence electrons are strongly attracted to the nucleus and an extra electron gives the atom a full valence shell. • Electron affinity decreases as we go down a group because the atoms are getting bigger and the valence electrons are not attracted as strongly to the nucleus.

25. Summary of Periodic Trends Nuclear charge increases Shielding is constant Atomic radius decreases Ionization energy increases Electronegativity increases Electron affinity increases Nuclear charge increases Shielding increases Atomic radius increases Ionic size increases Ionization energy decreases Electronegativity decreases Electron affinity decreases 1A 2A 0 4A 6A 3A 7A 5A Ionic size (cations) Ionic size (anions) decreases decreases