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Covalent Bonding and Naming

Covalent Bonding and Naming. I. Types of Covalent Bonds. l. Nonpolar covalent bond -a covalent bond in which the bonding electrons are shared equally 2. Polar covalent bond -a bond is which the bonded atoms do NOT share the bonding electrons equally. Types of Covalent Bonds.

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Covalent Bonding and Naming

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  1. Covalent Bonding and Naming

  2. I. Types of Covalent Bonds • l. Nonpolar covalent bond-a covalent bond in which the bonding electrons are shared equally • 2. Polar covalent bond-a bond is which the bonded atoms do NOT share the bonding electrons equally.

  3. Types of Covalent Bonds • If the difference in electronegativity between two bonded atoms is from 0.3 to 1.7, a polar bond will exist • If the difference in EN is less than 0.3 then the bond is nonpolar covalent.

  4. Practice • H and Cl • F and Br • S and I • O and H

  5. The Octet Rule • chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level.

  6. Multiple Bonding • single bond-two atoms share one pair of electrons • double bond-two atoms share two pair of electrons • triple bond-two atoms share three pair of electrons

  7. Drawing Lewis Dot Structures Add up the TOTAL number of valence electrons in the substance Decide what is the central atom. The central atom is the one that is least represented. (or the least electronegative) Arrange the dots so that each atom has an octet

  8. Practice 1. Br2 2. CCl4 3. H2O 4. NH3 5. O2 6. N2

  9. Exceptions to the Octet Rule • Some atoms have less than an octet. Example: Hydrogen only needs 2 electrons surrounding it and boron only needs 6. • H2 • BF3

  10. Exceptions to the Octet Rule Some atoms have more than an octet (One reason is because of bonding d orbitals as well as s and p orbitals.) Example: Sulfur can have up to 12 electrons surrounding it. SF6

  11. Resonance • – a concept in which two or more Lewis structures for the samearrangement of atoms (resonance structures) are used to describe the bonding in a molecule or ion. To show resonance, a double-headed arrow is placed between a molecule’s resonance structures. Example: O3

  12. Molecular Shapes The Valence Shell Electron Pair Repulsion Theory (VSEPR) is used to predict the shape of molecules. The VSEPR theory states that the bonding and nonbonding pairs of electrons will arrange themselves so that the repulsive forces between them are at a minimum. Molecules will adjust their shapes so that the nonbonding electrons are as far apart as possible.

  13. Molecular Shapes • The number of shared and unshared pairs of electrons determines the shape of the molecule.

  14. Linear • The bond angle is 180º . • 2 atoms bonded to the central atom • 0 lone pairs • Ex: CO2

  15. Bent The bond angle is 90º 2 atoms bonded to the central atom 2 lone pairs Write the Lewis structure for SF2

  16. Trigonal Planar The bond angle is 120° 3 atoms bonded to the central atom 0 lone pairs Write the Lewis structure for BCl3

  17. Trigonal Pyramidal The bond angle is 107° 3 atoms bonded to the central atom 1 lone pairs Write the Lewis structure for NH3

  18. Tetrahedral The bond angle is 109.5° 4 atoms bonded to the central atom 0 lone pairs Write the Lewis structure for CH4

  19. Trigonal bipyramidal The bond angle is 120° 5 atoms bonded to the central atom 0 lone pairs Write the Lewis structure for PCl5

  20. Octahedral The bond angle is 120° 6 atoms bonded to the central atom 0 lone pairs Write the Lewis structure for SF6

  21. Polarity A. nonpolar covalent bonds – EN difference is less than 0.3 There is equal sharing of electrons B. polar covalent bonds – EN is between 0.3 and 1.7 There is unequal sharing of electrons

  22. Practice • Determine whether each of the following is a polar covalent bond, nonpolar covalent, or ionic bond. • 1. N-O bond 2. Cl-Cl bond • 3. C-S bond 4. S-O bond • 5. P-O bond 6. Na-Cl bond

  23. Determining Polarity In a polar molecule there is an uneven distribution of charge. One end of the molecule is more negative or positive than the other In a nonpolar molecule there is an even distribution of charge.

  24. Determining Polarity Molecules with nonbonding pairs of electrons on the central atom are polar. H2O NH3

  25. Determining Polarity • The polarity of molecules with no nonbonding pairs depends on the atoms bound to the central atom. • If the surrounding atoms are identical, the molecule is nonpolar. This is because the bond dipoles cancel. • Example: CH4 • When the atoms surrounding the central atom are different, the molecule is polar. • Example: BBrI2

  26. Practice SO2 H2S CO2

  27. Intermolecular Forces of Attraction (IMFs)

  28. Van der Waals forces The attractive forces between molecules are collectively referred toas van der Waals forces. 1. There are very strong intermolecular forces of attraction between polar molecules and weak intermolecular forces of attraction between nonpolar molecules. 2. Many physical properties of a substance, such as freezing point and melting point, depend on the strength of the intermolecular forces of attraction.

  29. Dipole-Dipole Forces • 1. Dipole-dipole forces exist between polar molecules. The negative dipole of one molecule attracts the positive dipole of another molecule. • Example: HCl • 2. Dipole-dipole forces are weaker than chemical bonds.

  30. Hydrogen Bonding • 1. Hydrogen bonding is a special type of dipole-dipole force. Since no electrons are shared or transferred, hydrogen bonding is not a chemical bond. • 2. Hydrogen bonding always involves molecules containing hydrogen that is chemically bonded to a highly electronegative atom of small atomic size, specifically nitrogen, oxygen or fluorine Example: H2O

  31. London Dispersion Forces • 1. London dispersion forces are most noticeable in nonpolar molecules • 2. London dispersion forces arise from the motion of electron clouds . From the probability distributions of orbitals, it is concluded that the electrons are evenly distributed around the nucleus. However, at any one instant, the electron cloud may become distorted as the electrons shift to an unequal distribution. It is during this instant that a molecule develops a temporary dipole.This temporary dipole introduces a similar response in neighboring molecules, thus producing a short-lived attraction between molecules

  32. Strength of Forces of Attraction from Weakest to Strongest • 1. Chemical Bonds • a. Nonpolar Covalent • b. Polar Covalent • c. Ionic • 2. Van der Waals Forces • a. London Dispersion Forces • b. Dipole-Dipole Forces • c. Hydrogen Bonding

  33. Comparing Boiling Points • 1. When molar masses are similar, substances with stronger intermolecular forces of attraction have higher boiling and freezing points. • 2. In order to compare boiling and freezing points of a substances with similar molar masses you must: • a. Identify the structural formula • b. Determine the polarity of the molecule. • c. Determine the dominant type of Van der Waals forces • 3. Which of the following molecules has the higher boiling point: CH4 or NH3? Explain your answer.

  34. Naming and Writing Covalent Compounds • Writing Formulas for Binary Molecular Compounds those containing2 nonmetals. Use the prefix naming system - know theses prefixes

  35. Naming and Writing Covalent Compounds • mono – one • di – two • tri – three • tetra – four • penta - five • hexa – six • hepta – seven • octa - eight • nona - nine • deca – ten

  36. Practice • nitrogen tetrasulfide • carbon dioxide • oxygen monofluoride • sulfur hexachloride • trioxygen decanitride • tetrafluorine monophosphide • hexafluorine nonasulfide • heptabromine octanitride

  37. Naming Binary Molecular Compounds • The less electronegative element is given first. • The second element is named by combining a prefix indicating the number of atoms contributed by the element to the root of the name of the second element and then adding –ide to the end.

  38. Practice • CCl4 _________________________ • NF3 _________________________ • PBr5 _________________________ SF6 _________________________ • SO3 _________________________ • PCl5 _________________________ • N2O _________________________ PF6 _________________________

  39. Naming Acids

  40. Binary or hydrohalic acids Name “hydro____ic acid” Ex: HF ____________________ HCl ____________________ HBr ____________________ HI ____________________ H2S ____________________

  41. Oxyacids – contain a polyatomic ion If the polyatomic ion ends “ate” the acid will end in “-ic” HNO3 ____________________ H3PO4 ____________________ H2SO4 ____________________ H2CO3 ____________________ HC2H3O2 ____________________

  42. Oxyacids – contain a polyatomic ion If the polyatomic ion ends “ite” the acid will end in “-ous” HNO2 ____________________ H3PO3 ____________________ H2SO3 ____________________

  43. Metallic Bonds • Each metal donates its valence electron(s) to form an electron cloud • This leaves positive particles which are "cemented" together with the negative electron cloud, often called a “sea of electrons.”

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