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Covalent / Molecular Bonding

Covalent / Molecular Bonding. Sharing PAIRS of e- between non-metals and H Groups 4A, 5A, 6A, 7A No transfer of e- (ionic) Ways to represent covalent bonds. Examples. Single Covalent Bond between 2 hydrogen atoms

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Covalent / Molecular Bonding

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  1. Covalent / Molecular Bonding • Sharing PAIRS of e- between non-metals and H • Groups 4A, 5A, 6A, 7A • No transfer of e- (ionic) • Ways to represent covalent bonds

  2. Examples Single Covalent Bond between 2 hydrogen atoms • Lewis Structure must account for all valence e- to make a Lewis dot structure • __ 1 atomic orbital from each __ atom combine to form 1 molecular orbital c. Molecular Formula d. Structural Formula

  3. Examples Single Covalent Bond between 2 chlorine atoms • Lewis Structure must account for all valence e- to make a Lewis dot structure • __ __ __ __ 1 atomic orbital from each __ __ __ __ atom combine to form 1 molecular orbital c. Molecular Formula d. Structural Formula

  4. Examples Double Covalent Bond between 2 oxygen atoms • Lewis Structure must account for all valence e- to make a Lewis dot structure • __ __ __ __ here, 2 atomic orbitals from each __ __ __ __ atom combine to form 2 molecular orbitals c. d.

  5. Examples Triple Covalent Bond between 2 nitrogen atoms • Lewis Structure must account for all valence e- to make a Lewis dot structure • __ __ __ __ here, 3 atomic orbitals from each __ __ __ __ atom combine to form 3 molecular orbitals c. d.

  6. Examples • For covalent compounds, central atom in the molecule is listed first or underlined. • COUNT ELECTRONS!! • Hints: • H – single bond only • O –double bond or two singles • N – triple, double, or 3 singles (adds up to 3) • C – single, double, or triple bonds (adds up to 4) • Halogens - single bond only!!!

  7. Examples Practice drawing Lewis Structures and other ways of representing covalent compounds: • NH3 • CH4 • SiCl4 • CO2

  8. Examples Resonance Structures • To resonate – • Ex. SO2 Count – Octet • Ex. NO3-1 Count

  9. Exceptions to the Octet Rule • Expanded octets – 10 or 12 e- around the central atom • Ex. PCl5 Count – Ex. SF4 Count - ___BP ___ LP ___ BP ___ LP • Ex. XeCl2 Count ___ BP ___LP

  10. Why do these central atoms allow for expansion? • D orbitals become available for bonding • Atoms in periods 1 & 2 will never expand, but they will sometimes have incomplete octets.

  11. Incomplete Octets • Ex. BeF2 Count – Ex. BF3 Count - ___BP ___ LP ___ BP ___ LP

  12. Intro to VSEPRValence Shell Electron Pair Repulsion • A theory used to explain the SHAPE of a molecule or ion, based on its • number of shared (bonded) e- pairs • number of unshared (lone) e- pairs • e- pair arrangement / geometry • From the 3 items above, the following info can be determined about a molecule / ion: • the angles between e- pairs (bonded to atoms and lone) • the molecular geometry, the arrangement of atoms around the central atom • a model (sketch) can be drawn to depict geometry

  13. Steps for determining a molecule’s geometryUse your VSEPR Chart • Draw a Lewis structure, account for all electrons • Count the total # of e- pairs (BP & LP) • Select the correct VSEPR notation • Describe the electron pair geometry • Sketch a model of the molecule, accounting for e-pair repulsion • Show bond angles • Determine the Molecular Geometry

  14. VSEPR Practice • Ammonia, NH3 1. 2. 3. 4. 5. 6. 7.

  15. VSEPR Practice • H2O 1. 2. 3. 4. 5. 6. 7.

  16. VSEPR Practice • IF5 1. 2. 3. 4. 5. 6. 7.

  17. VSEPR Practice • BeF2 1. 2. 3. 4. 5. 6. 7.

  18. VSEPR Practice • CO3-2 1. 2. 3. 4. 5. 6. 7.

  19. VSEPR Practice • GeI2 1. 2. 3. 4. 5. 6. 7.

  20. VSEPR Practice • PCl4-1 1. 2. 3. 4. 5. 6. 7.

  21. Polarity • Table of electronegativity needed • Bonded pairs of e- pulled between nuclei of atoms that share the e- • H:H H:Cl • Non-polar covalent bond • same atoms bonded together, usually • bonding e- shared equally (0 - 0.4 Eneg Differences)

  22. Polarity • Polar covalent bond • two different atoms bonded together • bonding e- are shared unequally • Eneg Differences (0.4 - 1.9) • in the bond, the atom with stronger e- attraction (greater electronegativity) is slightly more negative ( -), and the other is more positive ( +)

  23. Polarity NONPOLAR POLAR IONIC The polarity of a molecule depends on the polarity of the individual bonds as well as the symmetry of the molecule (VSEPR structure).

  24. Practice determining type of bond using electronegativity table Bond H-H H-Cl P-F K-Cl Enegs Diff Polarity

  25. A dipole moment is indicated with an arrow in the direction of the more electronegative element in a bond, or to indicate the overall polarity of a molecule. Ex. HCl CO A polar molecule always has polar bonds. Some non-polar molecules have polar bond.

  26. Molecular Polarity • H2O Lewis BP e-pr. Geom. VSEPR Not. LP Molecular Geometry: Bond Polarity O-H Symmetrical? Molecular Polarity?

  27. Molecular Polarity • XeCl4 Lewis BP e-pr. Geom. VSEPR Not. LP Molecular Geometry: Bond Polarity Xe--Cl Symmetrical? Molecular Polarity?

  28. Molecular Polarity • CO2 Lewis BP e-pr. Geom. VSEPR Not. LP Molecular Geometry: Bond Polarity: C--O Symmetrical? Molecular Polarity?

  29. Molecular Polarity • PH3 Lewis BP e-pr. Geom. VSEPR Not. LP Molecular Geometry: Bond Polarity: P--H Symmetrical? Molecular Polarity?

  30. Molecular Polarity • CCl4 Lewis BP e-pr. Geom. VSEPR Not. LP Molecular Geometry: Bond Polarity: C--Cl Symmetrical? Molecular Polarity?

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