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Periodicity

Periodicity. Glencoe Chapter 6. Development of the Modern Periodic Table. 1790s – 23 known elements By 1870s – 70 known elements 1864 - John Newlands proposed arrangements by mass and properties by octaves

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Periodicity

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  1. Periodicity Glencoe Chapter 6

  2. Developmentof theModern Periodic Table • 1790s – 23 known elements • By 1870s – 70 known elements • 1864 - John Newlands proposed arrangements by mass and properties by octaves • 1864 – Lothar Meyer proposes arrangements by mass and columns of properties—but doesn’t publish! • 1869 – Dmitri Mendeleev also proposed arrangements by mass and columns of properties—and announces! • 1913 – Henry Moseley proposed arrangements by atomic number. (periodic law)

  3. Newlands’ “octaves”

  4. Shortly after, his ideas were presented to the Russian Physico-chemical Society. They were read by Professor Menschutkin because Mendeleev was ill. His ideas were then published in the main German chemistry periodical of the time, Zeitschrift fϋr Chemie. The world’s first view of Mendeleev’s Periodic Table – an extract from Zeitschrift fϋr Chemie, 1869. Click here for a translation

  5. Key “landmarks”of themodern periodic table • Periods (horizontal) • Groups/families (vertical) • Representative elements • s & p block • Groups 1A – 8A • Groups 1,2,13,14,15,16,17,18 • Transition elements • d block (f block = “inner transition elements”) • Groups 1B – 8B • Groups 3 - 12

  6. Other notable classifications: • Metals • Alkali (group 1) • Alkaline (group 2) • Metalloids • Nonmetals • Halogens (group 17) • Noble gases (group 18)

  7. Organizing Elements by Electron Configuration • Valence Electrons • Atoms in the same group have similar chemical properties b/c they have the same # if valence electrons • Valence Electrons and Periods • Energy level of valence electrons indicates the period on the PT • Valence Electrons and Group Number • Group # = # valence electrons • Exception: Helium

  8. The s-, p-, d-, and f- block elements • s-block • p-block • d-block • f-block

  9. Periodic trends • Vary systemically • across a period (horizontally) • along a group (vertically)

  10. Atomic radii • Based on probability of electron cloud, therefore, defined by how closely an atom lies to a neighboring atom • DECREASES to the right across a period • Due to larger nuclear attraction • INCREASES down a group • Due to more “layers” of electrons

  11. Ionic Radii • Ions (charged atoms) form when electrons are gained or lost….(the number of protons and electrons don’t match!) • DECREASES to the right across a period (in two phases) • INCREASES down a group

  12. Ionization Energy • Defined as “amount of energy required to remove an electron from a gaseous atom” • 1st ionization energy • 2nd ionization energy • Etc. • Think of this as the atom’s ability to hold onto its valence electron! • INCREASES across a period • Harder to remove e- • Positive energy means “harder” • DECREASES down a group • Easier to remove e- • More negative energy means “easier” or “more stable”

  13. Octet Rule • Atoms tend to gain, lose, or share electrons in order to acquire a full set of eight valence electrons • Note chemical stability of noble gases • Predicts ionic charge of main block elements • CATIONS—positively charged ions (lost e-) • ANIONS—negatively charged ions (gained e-)

  14. Electron Affinity • Energy associated with adding an electron to an atom’s electron cloud---think of the opposite of Ionization Energy…but same effect! • INCREASES (but the energy gets more negative = means “more stable”) across period • DECREASES (but the energy value gets more positive = means “more difficult”) down group • Therefore, a great idea!....

  15. Electronegativity • Indication of the relative ability of the atom to attract electrons in a chemical bond • Think of this quantity as how strongly an atom might want to gain an electron. • Arbitrary rating scaled to 4.0….. • Most electronegative element is fluorine with 3.98 • INCREASES across a period • DECREASES down a group

  16. Summary of Trends

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