Chemistry - Covalent Bonding
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Chemistry - Covalent Bonding

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Chemistry - Covalent Bonding Presentation Transcript

  • 1. Covalent Bonding
  • 2. Sections 8.1, 8.2, 9.3, and *8.4* Sections 6.1 16.1, 6.5 and *16.3*
  • 3. Remember…
    • Ionic bonds form between…
    • An ionic bond happens when one atom… and the other atom…
  • 4. Remember…
    • The definition of ionic bond is…
    • The chemical formula of an ionic compound represents a…
  • 5.
    • The four sentences above represent four of the most essential differences between molecular and ionic compounds.
    • These differences are so important that…
  • 6. Essay Question: Test 3-1
    • Define ionic bond and covalent bond. Outline and define, in detail, four major differences between ionic compounds and molecular (covalent) compounds.
  • 7. Molecular Compounds
    • Formed by covalent bonds
    • Ionic compounds are generally crystalline solids at room temperature.
    • Molecular compounds (CO 2 and water, for example, have VERY different properties.)
  • 8. Molecular Compounds
    • Molecular compounds are formed through covalent bonds.
    • Covalent bonds are created when atoms SHARE electrons, instead of gaining and losing them.
  • 9. Vocabulary
    • Molecule : group of atoms joined by covalent bonds
    • Diatomic molecules : molecules consisting of two atoms
    • Molecular formula : shows how many atoms of each element a molecule contains
  • 10. Think About It…
    • Chlorine is a diatomic element, meaning that it exists in its atomic state as two bonded atoms.
    • Draw two chlorine atoms.
    • Is the bond between these two atoms ionic or covalent? How do you know?
  • 11. Properties of Molecular Compounds
    • Covalent bonds usually occur between…
    • Often are gases or liquids at room temperature
    • Images will show atoms “stuck” to one another
  • 12. Properties of Molecular Compounds
    • In general, melting and boiling points of molecular compounds are lower than ionic compounds
  • 13. Molecular Formulas
    • Molecular formula of a molecular compound shows how many atoms of each element are in ONE MOLECULE of the compound.
    • (Contrast this with the chemical formula of ionic compounds, which show only the ratio of elements in the compound.)
  • 14. Molecular Formulas
    • Example:
    • IONIC: Calcium chloride – CaCl 2
      • Means that in the compound there are two chloride ions for every one calcium ion
    • COVALENT: Carbon dioxide – CO 2
      • Means that each carbon dioxide molecule consists of one carbon atom bonded to two oxygen atoms
  • 15. Molecular Compounds
    • Molecular compounds can be significantly larger than ionic compounds.
    • Benzoic Acid: C 7 H 6 O 2
    • 2,4-Dichlorophenoxyacetic acid: C 8 H 6 Cl 2 O 3
  • 16. Molecular Compounds
    • Formulas not always in lowest terms
    • Example: Ethane C 2 H 6
    • Formulas do not give molecule’s structure. (It must be inferred.)
  • 17. Structure Diagrams
    • Molecular Formula
    • Structural Formula
    • Ball-and-stick model
    • Space Filling Model
    • Perspective drawing
  • 18. Forming Covalent Bonds
  • 19. Octet Rule
    • In covalent bonds, atoms share electrons so that they fill their valence levels
    • Usually 8 (but only 2 for hydrogen)
  • 20. Single Covalent Bonds
    • Atoms held together by sharing one pair of electrons are said to form a SINGLE COVALENT BOND
    • Each atom donates one electron to the bond
  • 21. Single Covalent Bonds Cl Cl Cl Cl
  • 22. Single Covalent Bonds Cl Cl Cl Cl
  • 23. Single Covalent Bonds Cl Cl Cl Cl Single Bond Lone Pairs
  • 24. Covalent Bonds
    • Electrons that do not take part in the bond are called “lone pairs” or “unshared pairs”
  • 25. Covalent Bonds
    • Different elements can form different numbers of bonds
    • Group 7A elements need one more electron, and can form one bond
    • Group 6A elements need two more electrons and can form two bonds
    • Group 5A – three bonds
    • Group 4A – four bonds
    There are exceptions! Hydrogen, too!
  • 26. Working With Covalent Bonds
    • Draw the electron dot structures.
    • Determine arrangement.
    • Replace shared pairs of electrons with a line. (Leave lone pairs.)
  • 27. Draw Structural Formulas
    • NH 3
    • H 2 S
    • PBr 3
  • 28. Draw Structural Formulas
    • H 2 O
    • CH 4
    • OF 2
  • 29. Draw Structural Formulas
    • SCl 2
    • N 2 H 4
    • CCl 4
    • CHCl 3
    • C 2 H 6
    • HF
    Usually, the atom that can form MORE bonds will be in the center of the molecule!
  • 30. Draw Structural Formulas, Part 2
    • OBr 2
    • P 2 H 4
    • CI 4
    • CH 2 Br 2
    • C 2 Cl 6
    • HCl
    • C 3 H 8
    Usually, the atom that can form MORE bonds will be in the center of the molecule!
  • 31. Double and Triple Bonds
  • 32. Double Covalent Bonds
    • Atoms attain noble gas configuration by sharing two pairs of electrons (four)
    • Bond length is shorter
  • 33. Double Covalent Bonds
    • Oxygen has 6 valence electrons
    • O (Group 6A) can form two bonds
    O O
  • 34. Double Covalent Bonds
    • Oxygen has 6 valence electrons
    • O (Group 6A) can form two bonds
    O O
  • 35. Double Covalent Bonds
    • OCTET RULE NOT FULFILLED!
    O O
  • 36. Double Covalent Bonds
    • OCTET RULE NOT FULFILLED!
    O O
  • 37. Double Covalent Bonds
    • OCTET RULE FULFILLED!
    O O
  • 38. Other molecules with double covalent bonds are…
    • CO 2
    • Ethene, C 2 H 4
    • Carbonyl, COH 2
  • 39. Double Covalent Bonds
    • When counting number of valence electrons, double bonds count as 4 shared electrons.
    • Hydrogen will not form double covalent bonds… why?
  • 40. Triple Covalent Bonds
    • Atoms attain noble gas configuration by sharing three pairs of electrons (six)
    • Bond length is even shorter
  • 41. Triple Covalent Bonds
    • Nitrogen has 5 valence electrons
    • N (Group 5A) can form three bonds
    N N
  • 42. Triple Covalent Bonds
    • Nitrogen has 5 valence electrons
    • N (Group 5A) can form three bonds
    N N
  • 43. Triple Covalent Bonds
    • OCTET RULE NOT FULFILLED!
    N N
  • 44. Triple Covalent Bonds
    • OCTET RULE NOT FULFILLED!
    N N
  • 45. Triple Covalent Bonds
    • OCTET RULE NOT FULFILLED!
    N N
  • 46. Triple Covalent Bonds
    • OCTET RULE NOT FULFILLED!
    N N
  • 47. Triple Covalent Bonds
    • OCTET RULE FULFILLED!
    N N
  • 48. Other molecules with triple covalent bonds are…
    • Acetylene, C 2 H 2
    • Hydrogen Cyanide, HCN
    • Propyne, C 3 H 4
  • 49. Triple Covalent Bonds
    • When counting number of valence electrons, triple bonds count as 6 shared electrons.
  • 50. Diatomic Elements
  • 51. Diatomic Elements
    • Diatomic elements exist in their atomic forms as binary molecular compounds, since covalent bonds form between the atoms
    • i.e. a “molecule” of oxygen gas is O 2 , not O
  • 52. Diatomic Elements
    • F 2
    • Cl 2
    • Br 2
    • I 2
    • H 2
    • N 2
    • O 2
  • 53. Exceptions To The Octet Rule
  • 54. Exceptions
    • Compounds cannot satisfy the Octet Rule for all atoms if the total number of valence electrons is odd.
    • NO 2 – total number of valence electrons is 17
    O O N
  • 55. More Exceptions…
    • Nonmetals in the third period and beyond can form more than 4 bonds, since they have empty d orbitals where they can “promote” or “store” extra s or p electrons.
    • Ex.: Phosphorus can form 5 bonds.
  • 56. Naming Binary Molecular Compounds
  • 57. Naming Molecular Compounds
    • CO and CO 2 are very different compounds
    • How can we distinguish them in their names?
  • 58. Naming Molecular Compounds
    • Confirm that the compound is molecular, not ionic.
    • Name the elements in the order listed in the formula.
    • Add prefixes to identify the numbers of each atom in the compound.
  • 59. Prefixes Used Mono- 1 Di- 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6 Hepta- 7 Octa- 8 Nona- 9 Deca- 10
  • 60. Naming Molecular Compounds
    • Omit the prefix “mono-” on the first element in the name.
    • Add “-ide” as a suffix at the end of the second element’s name.
  • 61. Examples
    • N 2 O
    • Nitrogen oxygen
    • Dinitrogen monoxygen
    • DINITROGEN MONOXIDE
  • 62. Practice: Write the Molecular Formula
    • Nitrogen trichloride
    • Carbon tetrabromide
    • Diphosphorus trisulfide
  • 63. Practice: Write the Name
    • Cl 2 O 8
    • PH 3
    • N 2 O 4
    • SF 6
    • H 2 O
    • S 2 F 10
    • PCl 5
    • N 2 F 6
  • 64. Polar Bonds and Molecules
  • 65. Electronegativity
    • A measure of how well an atom attracts electrons
    • Measured in “Paulings”
    • In a molecule, some atoms more forcefully attract electrons than others
  • 66. Electronegativity
    • Decreases from top to bottom
    • Increases from left to right
  • 67. Polar Bonds
    • Polar bond – covalent bond in which electrons are shared UNEQUALLY
    • Difference in electronegativity values controls whether bond is nonpolar, polar, or ionic
  • 68. Polar Bonds
    • Differences:
    • 0.0-0.4  nonpolar covalent
    • 0.4-2.0  polar covalent
    • 2.0+  ionic
  • 69.  
  • 70. Polar Bonds
    • Greek letter Delta (δ) represents the partial charge acquired by atoms in a polar bond
    • H 2 O
    • HF
    • CO 2
  • 71. Intermolecular Attractions
    • Polar molecules attracted to one another (called dipole interactions)
    • Hydrogen bonds are attractions that occur between hydrogen and unshared electrons on another molecule
  • 72. Test Review
  • 73. Test Review
    • Covalent bonds – definitions
    • Molecular vs. ionic compounds
    • Writing structural formulas
    • Writing molecular formulas (from name or from structure)
    • Writing compound names
    • Information on Polar Bonds (pg. 237-240)