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GPS & Essential Questions

SC4 (a & b): Students will use the organization of the periodic table to predict properties of elements. (a): Use the Periodic table to predict trends including atomic radii, ionic radii, ionization energy, & electronegativity of various elements.

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GPS & Essential Questions

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  1. SC4 (a & b): Students will use the organization of the periodic table to predict properties of elements. (a): Use the Periodic table to predict trends including atomic radii, ionic radii, ionization energy, & electronegativity of various elements. EQ: What are periodic trends and how can they be used to characterize atoms and ions? GPS & Essential Questions

  2. Periodic Trends • In this chapter, we will rationalize observed trends in • Atomic Radii • Ionic Radii • Ionization energy • Electron affinity • Electronegativity

  3. Effective Nuclear Charge • In a many-electron atom, • electrons are both attracted to the nucleus • and repelled by other electrons. • The nuclear charge that an electron experiences depends on both factors.

  4. Effective Nuclear Charge The effective nuclear charge, Zeff, is found this way: Zeff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

  5. What Is the Size of an Atom? Atomic Radius: The bondingatomic radius is defined as one-half of the distance between covalently bonded nuclei.

  6. Atomic Radii: Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row (due to increasing Zeff). …increase from top to bottom of a column (due to increasing value of n).

  7. Ionic Radii: Sizes of Ions • Ionic size depends upon: • The nuclear charge. • The number of electrons. • The orbitals in which electrons reside.

  8. Ion Radii: Sizes of Ions • Cations are smaller than their parent atoms. • The outermost electron is removed and repulsions between electrons are reduced.

  9. Ionic Radii: Sizes of Ions • Anions are larger than their parent atoms. • Electrons are added and repulsions between electrons are increased.

  10. Ionic Radii: Sizes of Ions • Ions increase in size as you go down a column. • This is due to increasing number o energy levels.

  11. Ionic Radii: Sizes of Ions • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge.

  12. Ionization Energy • The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • The first ionization energy is that energy required to remove first electron. • The second ionization energy is that energy required to remove second electron, etc.

  13. Ionization Energy • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.

  14. Trends in First Ionization Energies • As one goes down a column, less energy is required to remove the first electron. • For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

  15. Trends in First Ionization Energies • Generally, as one goes across a row, it gets harder to remove an electron. • As you go from left to right, Zeff increases.

  16. Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e− Cl−

  17. Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row. Increases from bottom to top and left to right.

  18. The Periodic Table and Trends

  19. Valence Electrons: • Valence Electrons: Electrons in the outer-most energy level (Shell). • These electrons participate in bonding by losing, gaining, or sharing themselves. • 17 across the periods (Excluding the transitional metals) • Valence electrons hold atoms together in chemical compounds. • The negative charge of the valence electron is concentrated closer to one atom than to another. • This affects the chemical properties of a compound.

  20. Electronegativity: • Electronegativity: a measure of the ability of an atom in a chemical compound to attract electrons. • Period & Group Trends: • Most electronegative element is Fluorine • Electronegativities increase across each period, and decrease down a group or remain the same

  21. The periodic table of the elements in its modern form was first prepared by Dmitry Mendeleyev (1834-1907). Many of the chemical and physical properties of the elements change in a regular fashion, with successive elements in either the same row or the same column of this table. (A) How do the periodic trends change across rows and down columns in the table? Atomic radii, ionic radii, ionization energy, electron affinity, & electronegativity. (B) Explain, in terms of atomic structure, why each of these changes occurs in the manner observed. Checkpoint Question:

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