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Trends in Atomic Properties: Periodicity

Explore the general trends in the properties of atoms and their ions using the periodic table. This chapter discusses the relationship between atomic mass, atomic radius, and atomic number. Learn how atomic radius changes across periods and down groups, and how ionization energy is affected.

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Trends in Atomic Properties: Periodicity

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  1. Chapter 1 Activity 6: Atoms with More than one Electron

  2. PERIODICITY There are general trends in the properties of atoms and their ions. These trends can be explained using the periodic table.

  3. PERIODICITY Can you name a trend that we have talked about this year? Think about what you graphed • Atomic Mass vs Atomic # • As Atomic # increases, Atomic Mass increases • Atomic Radius vs Atomic # • As Atomic # increases, Atomic Radius increases down a group, but decreases across a period.

  4. Atomic Radius vs Atomic Number Period Trend As you proceed from left to right across a row (period), the atomic radius decreases (the atoms get smaller). Atomic Radius (pm) Why???? Atomic Number (protons)

  5. Notice the increasing number of protons…. Period The period # indicates how many energy levels the atoms in that period have. SO….as the number of protons increases, the outer electrons are pulled in closer to the nucleus and the atomic radius gets smaller!

  6. Atomic Radius vs Atomic Number Group Trend As you proceed down a group (family or column), the atomic radius increases (the atoms get larger). Atomic Radius (pm) Why??? Atomic Number (protons)

  7. The period # indicates how many energy levels the atoms in that period have. H Row 1 Li Row 2 So…As you proceed down a group, the number of energy levels increases. This increases the atomic radii of the atoms. Row 3 Na K Row 4

  8. http://www.youtube.com/watch?v=xjEJl88AYMo http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/atomic4.swf

  9. Assessment….. • Which atom is larger, Al or Si? Why???? • Which atom is smaller, Br or I? Why????

  10. Background Electrons can move between energy levels (Bohr’s model supported by line spectra) The nucleus and electrons are attracted to each other (opposites attract) Electrons that are far away from the nucleus can escape the attraction of the nucleus (atoms can lose electrons) Losing an electron is called ionization

  11. Ionization Energy Ionization energy is the energy required to remove a valence electron from an atom. Valence electron= outermost electron that is farthest from the nucleus. period group

  12. Ionization energy vs. atomic number

  13. Ionization Energy in kJ/mol D I

  14. Ionization Energy • Energy needed to remove a loosely held electron from an atom • Decreases as you move down a group • Atom gets bigger, more energy levels • Nucleus is further from the outer most electron (valence electrons) • Nucleus & electron attraction decreases, so it is easier to take an electron away. D

  15. Ionization Energy • Energy needed to remove a loosely held electron from an atom • Increases as you move across a period • Nonmetals hold onto their electrons tighter • Atom is smaller, so nucleus is closer to valence electron making it harder to remove. I • http://www.youtube.com/watch?v=o6tv03DR_Gg D

  16. Ionization Energy Practice He or Xe? (kJ/mol) Which has higher ionization energy? He Xe2372 kJ/mol 1170 kJ/mol Smaller atom, He holds on tighter to valence electron

  17. Ionization Energy Practice Li or Ne? (kJ/mol) Which has higher ionization energy? Li Ne520 kJ/mol 2081 kJ/mol Ne holds onto atoms tighter

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