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Chapter 11

Chapter 11 . The Periodic Table. I. History of the Periodic Table. Johann Wolfgang Döbereiner and triads John Newlands and the Law of Octaves Dmitri Mendeleev and the 1 st periodic table. Mendeleev’s Periodic Table. Mendeleev’s Predictions. Periodic Law.

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Chapter 11

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  1. Chapter 11 The Periodic Table

  2. I. History of the Periodic Table • Johann Wolfgang Döbereiner and triads • John Newlands and the Law of Octaves • Dmitri Mendeleev and the 1st periodic table

  3. Mendeleev’s Periodic Table

  4. Mendeleev’s Predictions

  5. Periodic Law • Basis: Element arranged according to their atomic masses present a clear periodicity of properties • Modern: The properties of elements repeat periodically when the elements are arranged in increasing order by their atomic numbers

  6. Circular Periodic Table

  7. Benfey’s Periodic Table

  8. Physics Periodic Table

  9. ADOMAH periodic table - electron orbitals

  10. Spiral Periodic Table

  11. We like spirals!

  12. Periodic system Pyramid format

  13. Periodic system: Zmaczynski & Bayley

  14. Periodic table in binary electron shells layout, designed by Eric William McPherson

  15. Regions of the Periodic table

  16. Representative Elements -EC • Valence v. core electrons

  17. Representative Elements - Ions • Generalization of atom/ion stability • Usually means 8 valence = octet rule

  18. Transition Elements - EC • Remember the exceptions to filling d orbitals

  19. Periodic Trends – Atomic Radii • Worksheet: Atomic Size • Why does atomic radius decrease across a period? • Higher # = more protons = higher core charge • Increased attraction between p+ & e- • e- pulled closer to nucleus = ???? • Why does atomic radius increase down a group? • Valence electron shell  higher n = higher probability of finding e- further from nucleus = ???? • Shielding by core e- = less pull on valence e- = ???? Smaller radius Larger radius Larger radius

  20. Atomic Radii

  21. Periodic Trends – Ionic Radii • Cation (+) radii are smaller than atomic radii • Why? • Lose of valence e- • Results in lower n, resulting in stronger nuclear pull • Anion (-) radii are larger than atomic radii • Why? • Gain of e- • Results in increased repulsion between e-

  22. Sizes of Anions (- ions)

  23. Sizes of Cations (+ ions)

  24. Graph of Atomic Radii

  25. Definition of Ionization Energy (IE) • Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The first or initial ionization energy or Ei of an atom or molecule is the energy required to remove one mole of electrons from one mole of isolated gaseous atoms or ions. You may think of ionization energy as a measure of the difficulty of removing electron or the strength by which an electron is bound. The higher the ionization energy, the more difficult it is to remove an electron. Therefore, ionization energy is an indicator of reactivity.

  26. Periodic trends – 1st Ionization Energy exceptions

  27. Periodic Trends – 2nd Ionization Energy *The teal colored cells represent ionization energies where the valence shell is now (n-1). (Why do you think there is such a large jump in the ionization energies when the n-1 shell is now valence?)

  28. Periodic Trends - Electronegativity • Definition • Increases across a period (L to R), decreases down a group (top to bottom)

  29. Electronegativity • Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. • The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to Cesium and Francium which are the least electronegative at 0.7.

  30. Electron Affinity!

  31. Electron Affinity Definition • the quantitative measure, usually given in electron-volts (eV), of the tendency of an atom or molecule to capture an electron and to form a negative ion.

  32. Periodic Trend for electron affinity

  33. Periodic Trends - All *Note: The electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion

  34. Wrap-up • Periodic Table Activity

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