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Covalent Bonding and Lewis Structures

Covalent Bonding and Lewis Structures. Ionic Bonding. Generally occurs between metals and nonmetals Can also occur with polyatomic ions. Ionic Bonding. Involves the transfer of electrons, followed by electrostatic attraction. Covalent Bonding. Generally occurs between nonmetals .

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Covalent Bonding and Lewis Structures

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  1. Covalent Bonding and Lewis Structures

  2. Ionic Bonding • Generally occurs between metals and nonmetals • Can also occur with polyatomic ions.

  3. Ionic Bonding • Involves the transfer of electrons, followed by electrostatic attraction.

  4. Covalent Bonding • Generally occurs between nonmetals. • Involvessharingof electrons, rather than transfer.

  5. Octet Rule • Atoms will acquire, through sharing or transfer, the electron configuration of a noble gas. • Most noble gases have 8 valence electrons • He is the exception

  6. Valence Electrons • The electrons in the highest occupied energy level How can you find the number of valence electrons for a representative element?

  7. Dot Models • The number of dots is equal to the number of valence electrons. P Example: Phosphorus

  8. Lewis Structure for Molecules • Each atom in the molecule is connected by bonds. • Bonds are shared pairs of electrons, and they are represented by a dash. • Pairs that are not shared are called unshared electrons.

  9. A single bond is created by one shared pair of electrons • A double bond is created by 2 shared pairs of electrons • A triple bond is created by 3 shared pairs of electrons.

  10. Drawing Lewis Structures Steps: • Count up total number of valence electrons in the molecule. Divide by 2. This is the number of pairs of electrons you must place. • In other words: # valence electrons ÷ 2 = # electron pairs

  11. Drawing Lewis Structures 2. Draw a skeleton of the molecule, joining atoms with single bonds. (1 shared pair) • Usually the first atom in the molecular formula is central. • Nature likes symmetry. • Hydrogen is never central! • Because he can’t have more than one electron pair!

  12. Drawing Lewis Structures 3. From the total pairs you counted in step 1, subtract the pairs you used in step 2. This will determine the number of pairs you have left to distribute.

  13. Drawing Lewis Structures 4. Distribute remaining pairs as unshared electrons around the atoms in the molecule. If you have too few pairs for each atom to have an octet you may need double or triple bonds somewhere in the molecule.

  14. Exceptions to the Octet Rule • Atoms with less than an octet.

  15. Exceptions to the Octet Rule • Atoms with more than an octet. Elements in period 3 or higher, such as sulfur, phosphorus and bromine are capable of holding up to 6 pairs (12 e-). This is called an expanded octet and should be used when you have extra pairs that won’t fit anywhere. Note – you cannot have an expanded octet and a multiple bond in the same molecule

  16. Exceptions to the Octet Rule • Molecules with an odd number of electrons. Total Number of Electrons = 5 + 6 = 11

  17. Molecular Shape • In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts whichever 3D geometry minimizesthis repulsion. • We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.

  18. Molecules with no unshared pairs around the central atom. These molecules are symmetrical!!!

  19. Molecules with unshared pairs around the central atom.

  20. VSEPR Shapes Linear Trigonal Bent Planar Triatomic Pyramidal Tetrahedral Trigonal Bipyramidal

  21. VSEPR Bond Angles 180° 120 ° <104.5 ° 107.5 ° 109.5 ° 90 °, 120 °

  22. VSEPR Content Frame

  23. Bond Polarity • In covalent bonds, shared pairs of electrons are pulled between the nuclei of atoms sharing them. • Sometimes electrons are pulled equally and sometimes they are not. • There are two types of covalent bonds: • Non-polar • Polar

  24. Nonpolar Covalent Bonds • The electrons are shared equally. • All diatomic elements are nonpolar.

  25. Polar Covalent Bonds • When the electrons are shared unequally. • Electronegativity is the ability to attract electrons. • The more electronegative atom will have a stronger attraction for the bonded electrons and will have a slightly negative charge. • The less electronegative atom will have a slightly positive charge.

  26. Example - HCl • Look up the electronegativity values for hydrogen and chlorine.

  27. Example: HCl Which is more electronegative? • H: 2.1 • Cl: 3.0 Chlorine has a slightly negative charge, while hydrogen has a slightly positive charge.

  28. There are two ways to communicate the polarity of HCl: H – Cl H – Cl  +  - The lowercase Greek letter delta shows that the atoms involved acquire only partial charges. The arrow points to the more electronegative atom.

  29. Example: Water Is hydrogen or oxygen more electronegative?

  30. Oxygen!!

  31. Example: Water Is hydrogen or oxygen more electronegative? Oxygen!  - O H H The O-H bonds are polar.  +  +

  32. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? • Br and Cl?

  33. 3.0 – 2.1 = .9

  34. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? • Br and Cl? Polar!

  35. 4.0 – 3.0 = 1.0

  36. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? • Br and Cl? Polar! Nonpolar!

  37. 3.5 – 1.0 = 2.5

  38. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? • Br and Cl? Polar! Nonpolar! Ionic!

  39. 3.0 – 2.8 = .2

  40. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? Ionic! • Br and Cl? Nonpolar! Polar! Nonpolar!

  41. Summary

  42. Molecule Polarity • If a molecule has all nonpolarbonds then the molecule is nonpolar. • If a molecule has a polar bond then the whole molecule is usually polar, but not always. • If the molecular geometry is symmetrical, the bond polarities cancel, and the molecule is nonpolar.

  43. Remember your SYMMETRICAL shapes! These molecules are symmetrical!!!

  44. Example: carbon dioxide CO2 has two polar bonds, but since the structure is linear (and symmetrical) the bonds cancel.  +  -  - O C O Nonpolar molecule!

  45. Example: methane CH4 has four polar bonds, but since the structure is tetrahedral (symmetrical) the bonds cancel.  + H  -  + H C H  + Nonpolar molecule! H  +

  46. Example: Water Water has a bent shape, due to unshared electrons (so NOT symmetrical)- The polarities do NOT cancel.  - Polar molecule O H H  +  +

  47. In a polar molecule, one end has a positive charge and the other has a negative charge. A molecule that has poles is called a dipolar molecule or a dipole.

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