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History of Periodic Table

History of Periodic Table. Chapter 5. History. 1860s – 60 elements discovered Cannizzaro - agreed on method to measure atomic mass Search for relationships between properties of elements. Dimitri Mendeleev. Organized elements by increasing atomic mass

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History of Periodic Table

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  1. History of Periodic Table Chapter 5

  2. History • 1860s – 60 elements discovered • Cannizzaro - agreed on method to measure atomic mass • Search for relationships between properties of elements

  3. Dimitri Mendeleev • Organized elements by increasing atomic mass • Noticed chemical and physical properties followed trend, or pattern  Periodic

  4. Mendeleev’s Table

  5. Henry Moseley • Worked with Rutherford looking at line-spectras • Noticed better pattern when elements were organized by increasing atomic # • Periodic Law: the physical and chemical properties of elements are periodic functions of their atomic #s

  6. Regions of Periodic Table Group Project

  7. Main Group Elements s and p block elements

  8. Group 1A are the alkali metals • Group 2A are the alkaline earth metals

  9. Group 7A is called the Halogens • Group 8A are the noble gases

  10. The group B are called the transition metals

  11. Top: Lanthanide Series Bottom: Actinide Series

  12. Periodic Properties

  13. Atomic Radii (Atomic Size) • Def: half the distance between the nuclei of identical atoms that are bonded together } Radius

  14. Atomic Radii - Group trends H Li • As we go down a group • Another energy level… • So the atoms get bigger. Na K Rb

  15. Atomic Radii - Periodic Trends • Go across a period the radius gets smaller. • Same energy level. • More nuclear charge. • Outermost electrons pulled in closer Na Mg Al Si P S Cl Ar

  16. Ionization Energy (IE) • An e- can be removed from any atom if there is enough energy A + energy  A+ + e- • Ion: atom or group of bonded atoms that has a (+) or(-) charge • Process that results in ion formed is ionization

  17. Valence Electrons • Def: The e- available to be lost, gained or shared to form chemical compounds • e- found in the outermost s and p sublevels

  18. Ionization Energy (IE) • Def: the energy needed to remove one e- from an atom (IE1 – first IE) • Atoms with HIGH ionization energy  hold on tight to their electrons

  19. IE – Group Trends • As you go down a group IE decreases • Electron further away from nucleus • Less attraction to nucleus, easier to take e-

  20. IE – Periodic Trends • IE generally increases from left to right • Increasing nuclear charge • More nuclear charge holds on tight to e- • Exact opposite of atomic radius

  21. IE2and IE3 • Energy required to remove additional e- • Energies keep getting higher and higher • e- that are left are being held closer to nucleus  harder to remove • Pg. 155

  22. Ionic Radii (Ionic Size) • Cation: positive ion • Always smaller than atom • Lost e-, now nucleus pulling in more on remaining e-s • Anion: negative ion • Always bigger than atom • Gaining e-, now e- are crowded and spread out (repulsion of like charges)

  23. Ionic Radii – Group Trends • Same as Atomic Radii • More energy levels as go down  size increases

  24. Ionic Radii – Periodic Trends • 2 sections • Metals on LEFT make CATIONs • Nonmetals on RIGHT make ANIONS • Cations (1A – 4A) Anions (5A – 8A) • Decrease as go across (L-R) due to increase nuclear charge

  25. Electronegativity • Valance e- are involved in forming bonds • Some atoms in a chemical bond attract the valance e- more than the other (tug of war) • Linus Pauling – electronegativity – measure of the ability of an atom in a chemical compound to attract e- from another atom in the compound

  26. Electronegativity – Group Trends • Tend to decrease down a group or remain about the same • Noble gases are NOT assigned electronegativities

  27. Electronegativity – Periodic Trends • Tend to increase as you go across the table • F – most electronegative • Fr – least electronegative

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