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Section 6.3

Explore the periodic trends of atomic radius, ionization energy, and electronegativity in this informative guide. Learn how these properties change across periods and groups, and understand the factors that influence these trends.

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Section 6.3

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  1. Section 6.3 Periodic Trends

  2. Objectives • Compare period and group trends of several properties • Relate period and group trends in atomic radii to electron configuration

  3. Atomic Radius • Electron cloud no defined edges • 90% probability of finding an electron • Size varies from substance to substance

  4. Radius: Metals ½ distance between adjacent nuclei in a crystal of the element

  5. Radius: Nonmetals • ½ distance between nuclei of identical atoms that are chemically bonded

  6. Trends within Periods • The decrease from left to right • All in same principal energy level • Increased nuclear charge as you move to the right, which draws electrons closer to the nucleus.

  7. Trends within Groups • Increase from top to bottom • Nuclear charge increases • Electrons are added to higher principal energy levels • Increase number of shells

  8. Periodic Trends: Atomic Radius DECREASES I N C R E A S E S

  9. Ions • Atoms that gain or lose electrons • negative or positive charge

  10. Ionic Radius • + charge = smaller radius • - charge= larger radius

  11. Lose Electrons/ +

  12. Gain Electrons/ -

  13. Trends within Periods • L to R: positive ions’ radii decreases • 5A: Size of larger negative ions decreases

  14. Trends within Groups • Ionic radii of both positive and negative ions increase as you move down a group

  15. Ionization Energy • Energy required to remove an electron from a gaseous atom • Energy required to remove first electron is the first ionization energy • Nuclei’s strength of hold on valence electrons

  16. Ions and Ionization Energy • 1A has low IE form positive ions • 8A has high IE do not form ions

  17. Beyond the 1st IE • 2nd ionization energy = amount of energy needed to remove an electron from a 1+ ion • 3rd IE = to remove e- from a 2+ ion

  18. Trends within Periods • First ionization energies increase L to R • Inc. nuclear charge= inc hold on e-

  19. Trends within groups • First ionization energies decrease from top to bottom • Valence electrons are farther from nucleus

  20. Periodic Trends: Ionization Energy INCREASES D E C R E A S E S

  21. Octet rule • Atoms tend to gain lose or share electrons to acquire a full set of 8 valence electrons • Used to determine type of ion an element is likely to form

  22. Electronegativity • Relative ability of an atom to attract electrons in a chemical bond. • Arbitrary units called Paulings

  23. Electronegativity Trends • Increases from L to R • Except Noble Gases • REASON: nuclear charge increases, more attraction to electrons in its outermost energy level • Increases from top to bottom • REASON: electrons are further away from the nucleus and better shielded from the nuclear charge

  24. Periodic Trends: Electronegativity INCREASES D E C R E A S E S

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