Periodic Table


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Periodic Table

  2. 2. The development<br />
  3. 3. J.W. Dobereiner<br />Classified elements into several sets of triads<br />TRIADS<br />Li, Na, K<br />Ca, Sr, Ba<br />Have similar chemical properties<br />Properties of the middle element are approximate averages of the first and third elements<br />
  4. 4. J.A.R Newlands<br />62 elements are already known<br />“law of octaves” (increasing atomic mass)<br />Similar properties: 8th and 1st , 9th and 2nd , 10th and 3rd<br />
  5. 5. Dmitri Mendeleev<br />Together with Lothar Meyer, they published nearly identical schemes for classifying the elements<br />Wrote the names of the elements nd its properties in cards and arranged the cards in various ways.<br />
  6. 6. He noticed that a periodic repetition of the properties of the elements could be observed when the elements were arranged in increasing atomic masses.<br />He eventually produced the first periodic table of elements.<br />
  7. 7. H.G.J. Moseley<br />Correctly hypothesized the fundamental property of each elements – the amount of positive charge in the nucleus = Atomic no.<br />Proposed that the correct way to arrange elements was with the increasing atomic number<br />
  8. 8. The PERIODIC LAW<br />Basis for the Periodic Table<br />When elements are arranged in increasing atomic number, heir physical and chemical properties show a periodic pattern.<br />
  9. 9. Reading the periodic table<br />
  10. 10. Groups or Families<br />The arrangement of the elements in vertical columns<br />Each family has the similar properties<br />
  11. 11. Period <br />Horizontal rows in the periodic table<br />
  12. 12. Labeling and naming groups<br />
  13. 13.
  14. 14. Periodic Trends<br />
  15. 15. Metals and Non - Metals<br />
  16. 16. Effective Nuclear Charge<br />The measure of the attraction between the nucleus and the electron<br />Also defined by the equation<br />Zeff = Z – S<br />Wherein,<br />Z = number of proton in the nucleus<br />S = average number of electrons that are between the nucleus and the electron in question<br />
  17. 17.
  18. 18. Greater Zeff<br />The greater the attraction between the nucleus and the electron<br />The electron are drawn closer to the nucleus<br />The atomic size is reduced<br />
  19. 19. Greater Shielding constant<br />The lower the Zeff<br />The lesser the attraction between electrons and the nucleus<br />
  20. 20. Electrons in the inner shell (lower n values), effectively shields the electrons in the outer shell (higher n values)<br />However the electrons in the same shell do not effectively shield one another<br />Example: electrons in the fourth energy level<br />
  21. 21.
  22. 22. Trend?<br />Effective nuclear charge increases across any row in the periodic table.<br />Effective nuclear charge increases slightly moving down a family/group<br />
  23. 23. Atomic Size<br />Half of the internuclear distance between adjacent atoms (atomic radii)<br />Trend?<br />Across a period (L-R) = decreases<br />This is due to the fact that moving from left to right across a period, the atomic number increases while the shielding factor, does not significantly increase; therefore, the Zeff increases thus pulling the electrons towards the nucleus<br />
  24. 24. Down a group = increases<br />Down a group the principal quantum number of the outermost electron increases<br />
  25. 25.
  26. 26. Check-up<br />Arrange the following atoms in order of increasing size:<br />P, S, As, Se<br />2. Arrange the following atoms in order of decreasing atomic radius: Na, Be, Mg<br />
  27. 27. Ionic Size<br />An estimate of the size of an ion in a crystalline ionic compound<br />From the relationship between the nuclear size and the atomic size, the size of the ion relative to its parent atom can be predicted<br />
  28. 28. <ul><li>Cations are smaller than their parent atom.</li></ul>When cations form, electrons are removed from the outer level<br /><ul><li>Anions are larger than their parent atom</li></ul>When ions form, electrons are added to the outer level<br />The increase in repulsion causes the electrons to occupy more space<br />
  29. 29. Isoelectronic series<br />isoelectronic – ions having the same number of electrons<br />Size decreases as the nuclear charge (atomic number) increases.<br />
  30. 30. Check-up<br />Arrange the atoms and ions in order of decreasing size: Mg2+, Ca2+, Ca<br />Which of the following atoms and ions is the largest ?<br />S2-, S, O2-<br />Arrange the ions in order of decreasing size: <br />S2-, Cl-, K+, Ca2+<br />
  31. 31. Ionization Energy<br />The minimum energy required to remove an electron from the ground state of the isolated gaseous atom or ion. <br />First Ionization energy<br />Energy needed to remove the first electron from a neutral atom<br />Second Ionization energy<br />Energy needed to remove the second electron<br />The greater the ionization energy, the more difficult it is to remove an electron<br />
  32. 32.
  33. 33. Check-up<br />Referring to a periodic table, arrange the following atoms in order of increasing first ionization energy.<br />Ne, Na, P, Ar, K<br />
  34. 34. Electron Configuration of ions<br />When electrons are removed from an atom to form cation, they are always removed first from the orbitals with largest available principal quantum number, n.<br />Example :<br />Li  Li +<br />
  35. 35. When electrons are added to an atom to form anion, they are added to an empty or partially filled orbital<br />Example:<br />F  F-<br />
  36. 36. Check-up<br />Write the electron configurations for the <br />Ca2+ ion<br />Co3+ ion<br />S2- ion<br />
  37. 37. SW<br />
  38. 38. Electron Affinity<br />Energy change that occurs when an electron is added to an atom<br />Measures the attraction, or affinity, of the atom for the added electron<br />
  39. 39. Ionization energy vs Electron Affinity<br />Ionization energy<br />Measures the ease with which an atom loses an electron<br />Electron affinity<br />Measures the ease with which an atom gains an electron<br />
  40. 40.
  41. 41.
  42. 42.
  43. 43. Electronegativity<br />Ability of an atom to attract electrons o itself<br />
  44. 44.
  45. 45.
  46. 46.
  47. 47. Worksheet<br />
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