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Covalent Bonding

Covalent Bonding. Chapter 9 Honors Chemistry Glencoe. 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-)

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Covalent Bonding

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  1. 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. Designated with “δ”. • The side with the stronger pull will have a - dipole. The weaker pull will have a + dipole.

  7. 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

  8. 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.

  9. 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

  10. 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

  11. 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.

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

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

  14. 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

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

  16. 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

  17. 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.

  18. 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+

  19. 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

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

  21. 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-

  22. 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

  23. 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.

  24. 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.

  25. 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.

  26. 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?

  27. 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.

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

  29. 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!

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

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