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Chemistry Ch. 8, Covalent Bonding

Chemistry Ch. 8, Covalent Bonding. Covalent Compounds: Are joined together by SHARING electrons with each other. Covalent bonds  Hold them together. H 2 O. C 6 H 10 O 5 (Cellulose). Molecule  atoms joined together by covalent bonds.

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Chemistry Ch. 8, Covalent Bonding

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  1. Chemistry Ch. 8, Covalent Bonding

  2. Covalent Compounds: Are joined together by SHARING electrons with each other. • Covalent bonds Hold them together.

  3. H2O

  4. C6H10O5 (Cellulose)

  5. Molecule atoms joined together by covalent bonds. • The smallest representative particle of a given compound. • Diatomic Molecule TWO atoms “…”

  6. Show the naturally occurring diatomics on the periodic table. Makes a “number 7”.

  7. Covalent Compounds (AKA Molecular Compounds) have lower melting and boiling points than ionic compounds.

  8. Bp/Mp Why? NaCl vs. H2O

  9. Molecular Formula: The chemical formula for a covalent compound. • Like: H2O

  10. Representing Covalent Compounds: • Structural formulas

  11. Some good rules to remember for Structural Formulas: • H will ALWAYS have ONE bond. • For neutral atoms: • F has 1 bond total • O has 2 bonds total • N has 3 bonds • C has 4 bonds total

  12. Covalent compounds are formed by SHARING Valence electrons. • Look at H2.

  13. Dot structure, Structural formula, Electron Config, • F2

  14. Make a structural formula for water.

  15. Make a structural formula for methane.

  16. Double and triple bonds: • N2 • CO2

  17. Make up examples of other structural formulas and let students practice.

  18. Polyatomic Ions. Ions made of more than one atom. • All the atoms TOGETHER have a net charge. • Polyatomic Ions like to stick together!

  19. Formal Charge: The charge that an atom has in a compound. • How to find it: take the number of valence electrons that the atom normally has – the number of “sticks” connected to the atom –the number of valence electrons shown.

  20. NH4+ • PO43-

  21. Optional to show this or maybe come back to it: • For + ions, Figure out how many total valence electrons there are, then do the structural formula. • For – ions, add the extra electrons at the end.

  22. HW: Chapter 8: 1-8, Also draw structural formulas for NF3, SBr2, C2H4, N2H2

  23. Drawing Activity: • Draw out a molecule of NaCl and a molecule of water. SHOW/EXPLAIN: • -How atoms are bonded together in each. • Which one is covalent (how do you know?) Which one is Ionic (how do you know?) • Also, draw 1. a crystal of NaCl and 2. a few water molecules. • -Explain how interactions for NaCl and for water lead to high melting points for NaCl and low for water.

  24. Look at Chlorate anion, page 225 (Teacher’s edition) ClO3-

  25. How to tell if a structural formula is legitimate: • Must show the correct total number of valence electrons. • Atoms like, if possible, to obey the octet rule • Negative charges, if present, should probably be on electronegative atoms. • Positive charges, if present, should probably be on atoms with low electronegativities. • Total charge must be correct

  26. Look at some other polyatomic ions on page http://www2.pvc.maricopa.edu/tutor/chem/chem130/nomenclature/polyatomicion.html • a) SO42- b) NH4+ c) CO32-

  27. Bond Dissociation Energy: the amount of energy it takes to break a chemical bond.

  28. C-C = 347 kJ/mol • C=C = 657 kJ/mol • H-H = 435 kJ/mol • C-H = 413 kJ/mol

  29. Resonance: Two or more different structural formulas are possible for the same compound. • Use a double arrow to show the possibilities.

  30. The actual bonding in a compound that has resonance is a Hybrid between the two forms shown.

  31. Sometimes there are exceptions to the octet rule, for example when a compound has an odd # total of valence electrons.

  32. SF6 has over an octet

  33. For Lab: Red = O Black = C Gray = H

  34. Why do atoms like to get an octet? • So they can achieve an electron configuration like a noble gas.

  35. Molecular Orbitals: • This theory states that when covalent bonds form, orbitals overlap in order to form a molecular orbital, which is shared by the whole compound.

  36. Sigma Bond • Two orbitals combine in order to form a symmetrical molecular orbital.

  37. Sigma bond

  38. Sigma bond • 2 p orbitals combine.

  39. Pi Bonds form when orbitals overlap side by side.

  40. P orbitals making a pi bond.

  41. Pi bonds only exist in compounds that have double or triple bonds.

  42. VSEPR Theory Explains the actual shape that compounds have.

  43. Unshared electrons strongly push away bonds.

  44. Methane (109.5) vs. ammonia (107)

  45. Why is CO2 linear but water is bent?

  46. Common Molecular Shapes, Fig. 8.18 • Tetrahedral • Trigonal Planar • Linear Triatomic • Bent Triatomic • Pyramidal • Trigonal bipyramidal • Octahedral • Square Planar • T-Shaped

  47. DISCUSS WHY THESE SHAPES ARE SEEN!!! • Tetrahedral • Trigonal Planar • Linear Triatomic • Bent Triatomic • Pyramidal • Trigonal bipyramidal • Octahedral • Square Planar • T-Shaped

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