1. Chapter 9 Honors Chemistry Glencoe
Covalent Bonding
Chapter 9
Honors Chemistry
Glencoe

2. Section 9-1: The Covalent Bond
Atoms bond so they can achieve a stable valence e- configuration (8e-, except for H & He, which need only 2e-)
Lewis Dot Structures: element surrounded by dots indicating the number of valence electrons. Examples.

3. Covalent Bond: shared e- form a bond between nonmetals to form a molecule
Electronegativity and Covalent Bonding: Sharing: Equally or Unequally? (Oreos)

4. Nonpolar covalent bond: electrons are equally shared (equal or almost the same electronegativity or pull). Ex: Cl2, O2.

5. Polar covalent bond: electrons are shared unequally (electronegativities are different) Section 9-5

6. Dipoles – poles with + and – charges are formed when the e- are pulled
Dipoles – poles with + and – charges are formed when the e- are pulled. Designated with “δ”.
The side with the stronger pull will have a - dipole. The weaker pull will have a + dipole.

8. Electronegativity and Bond Types – find the difference (subtract) the electronegativities
Nonpolar Covalent bond: if the difference is below 0.5
Polar Covalent bond: if the difference is between 0.5 and 1.7
Ionic bond: if the difference is greater than 1.7

9. Section 9-2: Naming Molecular Compounds and Acids
Naming and Writing Formulas for Covalent Compounds
Naming covalent (molecular compounds)
Name the element in the same order that they appear in the formula.
Change the last syllable of the last element to –ide.

10. Add prefixes to each element’s name to indicate how many atoms of that element are present in the formula.
Mono- is optional, especially for the first element in the formula.
Ex: CO2, CO, BF3, NI7, P4O5, P8 Cl9, N6O, CaCl2

11. Writing formulas of covalent (molecular) compounds (nonmetals with nonmetals)
Write the symbol of the elements in the order they appear.
Give them the subscript as indicated by prefix.
Examples: carbon monoxide, dicarbon tetrafluoride, triphosphorus hexasulfide, nonacarbon heptaiodide, copper oxide

12. Acid Names and Formulas
Acids are combinations of the hydrogen ion (H+) and an anion (negatively charged ion).
Examples: H+ and SO42- form H2SO4
H+ and Cl- form HCl
H+ and PO43- form H3PO4
Naming of acids are based on the endings of the anion. There are 3 endings: -ide, -ite and -ate.

13. Anions with –ide ending are named hydro___ic acid. Examples:
Anions with –ide ending are named hydro___ic acid. Examples: H+ and Cl- form HCl hydrochloric acid H+ and Br- form HBr hydrobromic acid

14. Anions with –ite ending are named ___ous acid. Examples:
Anions with –ite ending are named ___ous acid. Examples: H+ and ClO2- form HClO2 chlorous acid H+ and PO33- form H3PO3 phosphorous acid

15. Anions with –ate ending are named ___ic acid. Examples:
Anions with –ate ending are named ___ic acid. Examples: H+ and NO3- form HNO3 nitric acid H+ and C2H3O2- form HC2H3O2 acetic acid H+ and CO32- form H2CO3 carbonic acid

16. Section 9-3: Molecular Structures
Lewis Dot Structures
Dots (or crosses) represent valence e-.
Lewis Structures represent covalently bonded molecules.

17. Bonded pairs: pairs of e- that are shared are shown as a line
Unshared or lone pairs – e- that are not shared are shown as a pair of dots

18. Drawing Lewis Structures for Molecules
Count the total number of valence e-.
Arrange the atoms. Carbon is usually in the center. Hydrogens and halogens are normally on the ends.
All elements have 8 e-, except for H and He, which have 2 e-. Share electrons to accomplish this. Place 8 around central atom to start. Hint: E- do NOT need to stay with initial element!
Draw lines to represent the shared bonds.

19. Check your drawing: Is the number of e- correct? Is everyone happy?
Examples:
HCl
PH3
CF4
H2O
Structures for ions: extra or missing electrons are accounted for; brackets with a charge is used. Ex: NH4+

20. Multiple Bonds: sometimes more than one pair of electrons are shared resulting in double and triple bonds. Ex: HCN, C2H4
Limitations of the Octet Rule
Atoms with less than an octet.
Central atoms can be stable with less than an octet. (Usually boron is the central atom.)
Example: BF3

22. Atoms with more than an octet.
Noble gases can have more than an octet on the central atom.
Examples: XeF4 and SF6

23. Equivalent dot structures
Resonance structures: different dot structures without changing the arrangement of the atoms. Only e- positions change.
Resonance is represented by double-ended arrows (↔).
Example: Carbonate ion, CO32-

24. Section 9-4: Molecular Shape
VSEPR (Valence Shell E- Pair Repulsion)
E- will determine the shape of the molecule.
The e- pairs will orient themselves so that they are as far away from each other as possible.
Shapes: See Handout

25. Molecular Shapes
Linear – atoms can be connected in a straight line. The central atom has no unshared pairs of electrons. Molecules with only 2 atoms will always be linear. Bond angle = 180°. Examples: O2, HCl, CO2.
Trigonal Planar – central atom is bonded to 3 other atoms with no unshared pairs of e-. Does not follow the Octet Rule. Boron is typically the central atom. Bond angle = 120°. Example: BCl3.

26. Tetrahedral – central atom is bonded to 4 other atoms
Tetrahedral – central atom is bonded to 4 other atoms. The valence e- are arranged equally in 3-d. Bond angle = 109.5°. Examples: CH4, CF4.
Pyramidal – central atom is bonded with 3 other atoms, but also has an unshared pair of valence e-. Bond angle = 107°. Examples: NH3, PCl3.

27. Bent – central atom is bonded to 2 other atoms and has 2 unshared pairs. Bond angle = 105°. Example: H2O.
Trigonal bipyramidal – 5 atoms attached to the central atoms which has no lone pairs. The bond angles are 90° and 120°. Example: PCl5.

28. Octahedral – 6 atoms around the central atom, which has no unshared pairs. The bond angles are 90°. Example: SF6.
Examples:
What is the shape of the molecule PI3?
What is the shape of HCN?
What is the shape of the molecule to the right?

29. Section 9-5: Polarity Molecular Polarity
Polar molecules – net pull in one direction; polarity is indicated with arrows. Be sure to use the Ball and Stick models! Example: HF, CHCl3 & H2O.

30. Non-polar molecules – overall nonpolar because all the pulls cancels each other out, but may still have individual polar bonds. Example: BF3 and CH4.

32. Solubility
“Like will dissolve like”: polar (ionic) will dissolve in polar and nonpolar will dissolve in nonpolar.
Water (polar) will not mix with oil (nonpolar)
Salt (ionic) will dissolve in water (polar), but salt will not dissolve in oil (nonpolar)…unless the temperature is changed!

33. Covalent Network Solid
Strong covalent bonds, therefore these solids have very high temperatures.
Examples: quartz, diamonds, asbestos, graphite