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Chemistry 104 Chapter Four

Chemistry 104 Chapter Four. Chemical Bonding: The Covalent Bond Model. You will be able to:. Compare and contrast ionic to covalent bonds Explain how atoms can have different numbers of covalent bonds Explain the term electronegativity and how it is used

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Chemistry 104 Chapter Four

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  1. Chemistry 104 Chapter Four Chemical Bonding: The Covalent Bond Model

  2. You will be able to: • Compare and contrast ionic to covalent bonds • Explain how atoms can have different numbers of covalent bonds • Explain the term electronegativity and how it is used • Explain the concept of polar molecules

  3. Covalent vs Ionic bonds • Ionic bonds are exchanges of electrons • Usually __________________combinations form binary ionic compounds

  4. Covalent vs Ionic bonds • Covalent bonds are a sharing of valence electrons to reach the octet number (8) • Occur btwn atoms that both want e-’s to ____________________________, so they share • ____________ and____________ combo’s tend to be covalently bonded.

  5. e- sharing can occur only when e-’s located btwn 2 nuclei are in a stable configuration btwn Covalent bonds can occur btwn atoms of the

  6. Formation of a Covalent Bond

  7. Examples H2 H : H F : F F2 H : Cl HCl

  8. # of covalent bonds formed - directly related to Oxygen, with 6 valence e-, forms ______ ___________ covalent bonds.

  9. Atoms needing multiple e-’s to fill valence shells can make multiple covalent bonds • Oxygen: Group VIA (or 16), has how many valence e-’s? • ______ valence e- & needs how many to get an octet? • _____ more to make ______________ • Oxygen makes ______ single covalent bonds can also make ______double covalent bond

  10. Double Bonds • O2 • C2H4 • CO2 • CH2CHCHCH2

  11. Nitrogen (N), Group VA (or 15) has 5 valence e-’s, forms three single covalent bonds. can also form ________triple bond or______ double and______ single bond.

  12. Triple bonds • N2 • C2H2 • HCN

  13. C, group IVA (14) has 4 valence electrons, forms _________ covalent bonds.

  14. Carbon “backbone” of most molecules in living organisms  bonding variability. C makes straight chains with itself, branched chains, rings, and groups of rings and chains

  15. Carbon can make:

  16. Carbon can make: 4 single bonds; 2 double bonds; 1 single and 1 triple bond; 1 double and 2 single bonds

  17. Vast majority of compounds contain these four elements: H, O, N, C Hydrogen = 1 single Oxygen = 2 singles or Nitrogen = 3 singles or Carbon = 4 singles or 2 doubles or

  18. Lewis Structures Structural drawing – shows • ________________ of atoms • ____________ of covalent bonds • SEVEN basic steps to follow in drawing Lewis Structural models.

  19. Step 1: Add up the # of valence electrons in the molecule H each have O has_______________ e.g. H2O

  20. Step 2: • Choose a “center” atom • Usually the single atom when several are present. • Central atoms: never CO2 H2O NH3 SCl2 PBr3 CCl4 NO2

  21. Draw single bonds (use dashes) from center atom to each other atom. • Try to make the drawing as symmetrical as possible. Cl – S - Cl

  22. Subtract # of electrons in bonds from

  23. Place remaining e-’s around atoms in pairs (use dots) until each atom in molecule obeys

  24. Step 6: If it appears there are not enough ELECTRONS Return to step 3, and substitute a double (or triple) bond for one of the single bonds.

  25. Return to step 5, and place the extra e-’s around the center atom in pairs. Methane CH4 Sulfur dioxide SO2 Ammonia NH3

  26. Nature never makes it easy for scientists to classify things CO2 is easy to see how the atoms are covalently bonded: O=C=O But what is up with SO2 (sulfur dioxide)? How is that bonding?

  27. (a) A “regular” covalent single bond is the overlap of two half filled orbitals. (b) A _________________covalent single bond is the result of overlap of a filled and a vacant orbital.

  28. How is sulfur dioxide bonding? • With coordinate covalent bonds! • Sulfur has ______ valence e- • Each oxygen has _______ valence e- • Total valence e-= • Write each symbol selecting central atoms 1st • Put in a single covalent bond (one pair of e-) • for each atom

  29. Next: • Add non-bonding e- pairs so that atoms bonded to the central atom have octets

  30. Next: Place remaining electrons on CENTRAL atom

  31. If more electrons are needed, use non-bonding (lone pair) electrons to create an octet # around the central atom.

  32. Resonance • Some molecules have measured values of bond lengths which do not support the Lewis structure drawn for the molecule • Example:

  33. Resonance • Example: Ozone, O3 • To adequately represent such molecules with Lewis structures, you must draw all possible arrangements of electrons.

  34. If you have electrons left over, the center atom does not obey the OCTET rule - it is an exception!How?

  35. Exceptions: Only elements in rows 3 and beyond. Why? • PCl3 and NCl3 • PCl5 but not NCl5 • KrF4 • SF6 Think about at what level the d subshell begins to fill…

  36. Linus Pauling Home grown Oregon boy does good Wins two Nobel Peace prizes! Develops numerical scale

  37. Electronegativity • The relative attraction an atom has for • shared electrons in a chemical bond. • Electronegativity_____________ • going • Electronegativity_____________ • going

  38. Abbreviated periodic table showing Pauling electronegativity values for selected elements. What’s the most electronegative element? Helps explains why it is so violently reactive.

  39. PolarvsNon-polar covalent bonds • Two atoms of different electronegativity will share their electrons unequally • This will produce regions of weak positive and negative fields or “_______” around the atoms • Use the Greek letter Delta with a superscript charge: e.g. HCN Hydrogen cyanide

  40. By looking at differencesbtwnelectronegativitiesof bonding atoms we can determine what type of bond will form:

  41. Non-polar covalent bonds must have __________ difference in electronegativity: • A difference > 2.0 will result in • Between those numbers will produce

  42. In symmetric molecules, no matter what the ________________ are, the net__________________of the molecule is zero because the bond dipoles cancel.

  43. If a molecule made up of polar bonds is ___________________then charges will not cancel out:the molecule is said to be “___________”

  44. Can you: • Compare and contrast ionic to covalent bonds • Explain how atoms can have different numbers of covalent bonds • Explain the term electronegativity and how it is used • Explain the concept of polar molecules

  45. Chemistry 104 Chapter Four Chemical Bonding: Covalent Bond Model Part 2 Tetraphosphorous decoxide Carbon tetrafluoride Dihydrogen dioxide

  46. You will be able to: • Review the term electronegativity and how it is used • Review the concept of polar molecules • Explain Valence Shell Electron Pair Repulsion theory (VSEPR) • Describe the geometry of covalent molecules using VSEPR • Name binarymolecular compounds

  47. Molecules have a 3-D structure formed from their electron arrangements • Electrons have similar negative charges • Similar charges __________ • opposite charges ____________ • e-’s orient themselves to avoid each other

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