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Ionic Compounds Containing Transition Metals (p. 55)

Ionic Compounds Containing Transition Metals (p. 55). Transition Metals: Includes Groups 3 to 11 The number of electrons that a transition metal atom can lose can vary. EX: Iron (Fe) atom can lose either 2 electrons OR 3 electrons forming Fe 2+ OR Fe 3+ respectively.

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Ionic Compounds Containing Transition Metals (p. 55)

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  1. Ionic Compounds Containing Transition Metals (p. 55) • Transition Metals: Includes Groups 3 to 11 • The number of electrons that a transition metal atom can lose can vary. • EX: Iron (Fe) atom can lose either 2 electrons OR 3 electrons forming Fe2+ OR Fe3+ respectively. • *****When you are working with transition metals, you will be given the charge or enough information to determine the charge on the ions

  2. Chemical Bonds (Molecular Compounds & Covalent Bonds) Part 2 Textbook Pages: 52-63

  3. Recall Definitions: • Covalent Bond: The attraction between atoms that results from the sharing of electrons. • Molecular Compounds: A chemical compound that is held together by covalent bonds. • Polar Covalent Bond: Electrons are shared unequally. ΔEN = 0.6 to 1.7 • Covalent Bond: Electrons are shared equally. ΔEN = Equal to or less than 0.5

  4. Formation of Covalent Bonds • Only unpaired electrons are likely to participate in chemical bonds. • How can you tell if electrons are shared in a bond? • You find the ΔEN!!! • Examples: • Carbon Dioxide: ΔEN = 1.0 (Polar Covalent) • Water: ΔEN = 1.4 (Polar Covalent)

  5. Rules for Drawing Lewis Structures for Molecular Compounds • 1. Determine, from the chemical formula, the number of atoms of each type of element in the compound. Use the last digit of the group number from the Periodic Table to determine the number of valence electrons for each atom. • 2. The element that requires the most electrons to complete its valence is called the central atom. Draw the Lewis structure for this atom by placing one electron on each side of the imaginary square enclosing the central atom before paring up electrons. • A) If there is more than one of this type of atom, write them side by side and bond them together using a line. Bonds form where there are unpaired electrons. • 3. Use the unpaired electrons to bond additional atoms with unpaired electrons to the central atom until a stable octet is obtained (there are some exceptions that disobey the octet rule).

  6. Drawing Lewis Structures for Molecular Compounds • Draw the Lewis structures for the following compounds: • H2 • Water (H2O) • Carbon Dioxide (CO2) • Nitrogen gas (N2) • OF2 • C2H2

  7. Atoms that can Disobey the Octet Rule • NOTE: There are some atoms that can disobey the octet rule. • This occurs when they are bonded to highly electronegative atoms (i.e., F, Cl, Br). • When this happens, they can take more than 8 electrons (i.e., P, N, S) or less than 8 electrons (i.e., Be, B).

  8. Drawing Lewis Structure for Molecular Compounds that Atoms that Disobey Octet Rule • Draw the Lewis structures for the following compounds: • BF3 • SF6 • PCl5 • BeF2

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