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Unit 6 Covalent Bonding

Unit 6 Covalent Bonding

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Unit 6 Covalent Bonding

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  1. Unit 6 Covalent Bonding Fructose Carbon Dioxide Ammonia

  2. 1/29/16 • Objective: Students will be able to draw Lewis dot structures for covalent compounds and differentiate between metallic, ionic and covalent bonds. • Due Today: Binders! • Agenda: • Warm-up • Notes/Practice • Binder Turn-in • Exit Ticket • Homework: • Covalent Bonding Practice Worksheet – Due 2/2 • Warm-up • What types of elements form ionic bonds? (Metals, non-metals, both?) • What types of elements form metallic bonds (Metals, non-metals, both?) • What types of elements form covalent bonds (Metals, non-metals, both?)

  3. Molecules and Molecular Compounds • Compounds that are NOT held together by an electrical attraction, but instead by a sharing of electrons. • Atoms held together by a sharing of electrons are joined by a covalent bond.

  4. H••H

  5. Molecules and Molecular Compounds • A molecule is a neutral group of atoms joined together by covalent bonds. A compound composed of molecules is called a molecular compound. • The chemical formula for a molecule is called the molecular formula.

  6. Properties of Molecular Compounds • Composed of two or more nonmetals. • Usually gases or liquids at room temperature.

  7. Properties of Molecular Compounds • Molecular compounds tend to have a lower melting and boiling point than ionic compounds. • Reason: There are no (or few and weak) bonds holding the molecules together in molecular compounds. Molecular Compound Ionic Compound

  8. Properties of Molecular Compounds • Do not conduct electricity. They form nonelectrolytes in solution. • Reason: Molecular compounds do not break apart into ions in solutions.

  9. Covalent Bonding and the Octet Rule • Compounds that are NOT held together by an electrical attraction, but instead by a sharing of electrons. • Atoms attain an octet (also called noble gas electron configurations) by sharing electrons. • The bonds that form from this sharing can be single, double or triple. Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.

  10. To draw Lewis structures for covalent bonds, use the NAS method: • N (Needed): Find the number of electrons needed to form full octets for all elements involved. For most nonmetals, the number needed is 8. Hydrogen is the exception, it needs only 2. • A (Available): Find the number of electrons available by adding up all of the valence electrons for all elements involved.

  11. NAS method continued: • S(Shared): Subtract the two numbers. S = N-A A bond is formed with two electrons, so divide by two to tell you how many bonds to draw between the elements. • Draw the molecule. Put first atom in the center. H’s are always outside. Draw in the bonds, then fill in the rest of the electrons. • Check to ensure all atoms have a full octet.

  12. Example #1: CH4 • N = • C needs 8 valence e- 1 C  1 x 8 = 8 • Each H needs 2 valence e- 4 H’s  2x4 = 8 • Add valence e- needed for C and H’s together 8+8 = 16 total • A = • C has 4 valence e-  4 • H has 1 valence e-  4 • Total Available = 8 • S = N-A = 16-8 = 8/2 = 4 bonds. • Draw • All Atoms have octet?

  13. Example #2: CO2 • N = • C needs 8 valence e- 1 C  1 x 8 = 8 • Each O needs 8 valence e- 2 C’s  2x8 = 16 • Add valence e- needed for C and H’s together 8+16 = 24 total • A = • C has 4 valence e-  4 • O has 6 valence e-  2x6  12 • Total Available = 16 • S = N-A = 24 - 16 = 8/2 = 4 bonds. • Draw • All Atoms have octet?

  14. Example #3: N2 • N = • N needs 8 valence e- 2 N  2 x 8 = 16 • A = • N has 5 valence e-  2 x 5  10 • Total Available = 10 • S = N-A = 16 - 10 = 6/2 = 3 bonds. • Draw • All Atoms have octet?