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The Periodic Table and Physical Properties

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  1. SONG The Periodic Table and Physical Properties Get a periodic table out. Topics 3.1 - 3.3 and 12.1.1 - 12.1.2

  2. Dmitri Mendeleev 8 February 1834 – 2 February 1907 • Russian chemist and teacher • given the elements he knew about, he organized a “Periodic Table” based on increasing atomic mass (it’s atomic # now) • he even left empty spaces to be filled in later (TOK– he was a “scientist” and “risk taker”!)

  3. At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic table.He predicted their discovery and estimated their properties.

  4. Z

  5. Design of the Table • Groups are the vertical columns. • elements have similar, but not identical, properties • most important property is that they have the same # of valence electrons

  6. valence electrons- electrons in the highest occupied energy level

  7. Electron arrangement (SL level – 3.1.3) 2 2,1 2,8 2,3 2,5 2,8,2 http://images.google.com/imgres?imgurl=http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/imgper/econfig.gif&imgrefurl=http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/perlewis.html&h=267&w=512&sz=22&tbnid=__EXctBwlG0J:&tbnh=66&tbnw=128&hl=en&start=1&prev=/images%3Fq%3DElectron%2BDot%2BDiagrams%26svnum%3D10%26hl%3Den%26lr%3D

  8. B is 1s2 2s2 2p1; • 2is the outermost energy level • it contains 3 valence electrons, 2 in the s and 1 in the p • Br is [Ar] 4s2 3d10 4p5How many valence electrons are present?

  9. Periods are the horizontal rows • do NOT have similar properties • however, there is a pattern to their properties as you move across the table that is visible when they react with other elements

  10. Definitions • atomic radii • the distance from the nucleus to the outermost electron • ionic radii • same distance, but for ions (atoms that have lost or gained valence electrons) • first ionization energy (kJ mol-1) • the energy needed to remove the outermost, or highest energy, electron from a neutral atom in the gaseous phase

  11. electronegativity • measures the attraction for a shared pair of electrons • melting point • chemical properties • how elements react with other elements

  12. Trends in the table

  13. But first, the electron shielding effect • electrons between the nucleus and the valence electrons repel each other

  14. H Li Na • Atomic radii • McGraw Hill video • groups (alkali metals and halogens) • increases downwards as more levels are added • periods across the periodic table (period 3) • radii decreases • the number of protons in the nucleus increases • increases the strength of the positive nucleus and pulls electrons closer to it K Rb

  15. + + Li , 0.078 nm 2e and 3 p • Ionic radii • decreases across periods for same reason as atomic radii (nucleus becomes stronger) • alkali metals • cations are smaller that the parent atom • have lost an electron (actually, lost an entire level) • therefore have fewer electrons than protons • radii still increases downwards as more levels are added on forming a cation Li 0.152 nm 3e and 3p

  16. - - F 0.064 nm F 0.133 nm 9e- and 9p+ 10 e- and 9 p+ • halogens • anions are larger than parent atom • have gained an electron to achieve noble gas configuration • radii still increases downwards as more levels are added on forming an anion

  17. Ionization energy • decreases down a group • outer electrons are farther from the nucleus and therefore easier to remove • inner core electrons “shield” the valence electrons from the pull of the positive nucleus

  18. increases across a period • extra electrons are just filling up the same level • the nucleus is becoming more powerful and therefore the electrostatic force increases making it harder to remove an electron

  19. 12.1.1 • Evidence for levels and sub-levels • First ionization energy • electrons are harder to remove… • when there are more protons to attract them • a sub-level (s,p,d,f) is completely filled • a sub-level (s,p,d,f) are half filled

  20. 12.1.1 • Evidence for levels and sub-levels • successive (1st, 2nd, 3rd) ionization energy • as more electrons are removed, the electrostatic pull of the protons holds the remaining electrons closer • therefore, more energy is required to remove them (even have to use a logarithmic scale to show this) • large “jumps” are when the electrons are being removed from the next, lower level that are much closer to the nucleus

  21. starting to remove the 2s sub-level starting to remove the 3s sub-level starting to remove the 1s sub-level starting to remove the 2p sub-level starting to remove the 3p sub-level 4s1 removed

  22. Electronegativity • as you go down a group electronegativity decreases • the size of the atom increases • the bonding pair of electrons (-) is increasingly distant from the attraction of the nucleus (+) • the valence electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electrons

  23. as you go across a period • electronegativity increases • the atoms become smaller so the positive nucleus can hold onto the electrons better

  24. Melting point • group 1 (alkali metals) • decreases as “sea of negative electrons” are farther away from the positive metal ions • group 7 (halogens) • increases downwards as the van der Waals’ forces increase • larger molecules have more electrons which increases the chance that one side of the molecule could be negative

  25. decreases increases

  26. across the table (period 3) • from left to right • bonding goes from strong metallic to very strong macromolecules (network covalent) to weak van der Waals’ attraction

  27. 1+ charge • Chemical properties • groups • alkali metals • react vigorously with water and air • 2Na (s) + H2O(l)  2Na (aq) + 2OH- (aq) + H2 (g) • (Li, Na, K… all the same equation) • reactivity increases downwards • because the outer (valence) electron is in higher energy levels (farther from the nucleus) and easier to remove • react with the halogens • halogens’ reactivity increases upwards • smaller size can attract electrons better • (see next slide) 1- charge

  28. least reactive most reactive

  29. halogens • diatomic molecules such as F2, Cl2, Br2, I2 • can react with halide ions (Cl-, Br -, and I -) • the single bond is broken and each atom can gain one electron to form halide ions (F1-, Cl1-, Br1-, I1-) • the most reactive ends up as an ion (1- charge) and is not visible (molecules F2, Cl2, Br2, I2 are a visible gas) • Cl> Br > I

  30. periods • from left to right in period 3 • metals…metaloids…nonmetals • when oxides react with water • basic…amphoteric (either basic or acidic)…acidic • Na2O(s) + H2O (l)  2 NaOH(aq) strong base • MgO (s) +H2O (l)  Mg(OH)2 (aq) weaker base • P4O10 (s) + 6H2O (l)  4 H3PO4 (aq) weak/strong acid • SO3(g) + H2O (l)  H2SO4 (aq) strong acid