Chemical Bonding

1. Chemical Bonding 1

2. Video http://ed.ted.com/lessons/how-atoms-bond-george-zaidan-and-charles-morton

3. Chemical Bonding • Atoms lose, gain, or share electrons to have an octet (8e-) in its outer energy level.

4. Types of Bonding • Ionic • Covalent • Metallic

5. Ionic Bonding • Transfer of electrons • Metal (loses e-) forms a + ion nonmetal (gains e-) forms a negative ion • Ex: NaCl, KBr, KNO3

7. e– 1) 2) Na Cl Cl– Na+ Ionic Bonding Example: Na and Cl In ionic bonding one atom has a stronger attraction for electrons than the other, and “steals” an electron from a second atom 3)

8. Properties of Ionic Compounds • Crystalline solids • High melting point • Brittle • Conducts electricity in molten form and when dissolved in water • Forms a crystal lattice • Held together by electrostatic attraction (opposite charges)

9. Crystal Lattice

10. Sodium and Chlorine NaCl

11. Calcium and Chlorine CaCl2

12. Potassium and Sulfur K2S

13. Covalent (Molecular) Bonding • Share e- • Nonmetals only • Ex: H2O, CCl4

15. Electronegativity • Measure of the attraction for a shared electron. • Which family would have the highest electronegativity? • Which element would have the highest electronegativity?

16. Electronegativity increases decreases Don’t include the noble gases

17. Electronegativities and Bond Type The type of bond or degree of polarity can be calculated by finding the difference in electronegativity of the two atoms that form the bond.

18. Two types of covalent bonds • Nonpolar bond- share e- equally • Ex: Cl2, O2 • Polar bond- unequal sharing of e-, 2 different nonmetals • Ex: C-O, N-H

19. Electronegativity Difference • Used to determine if a bond is ionic or covalent • Atoms that have electronegativity differencesgreater than 1.7 form ionic bonds. Ex: Na-Cl • If thedifference is 0 to 0.4 the bond is nonpolar covalent.Ex: Cl2. • Atoms that have electronegativity differencesless than 1.7 and greater than 0.4 form polar covalent bonds. Ex: H-O • The greater the electronegativity difference the more polar the bond will be. 19

20. Electronegativity Difference 0------------0.4-------------1.7------------ nonpolarpolar ionic

21. Examples Classify as polar, nonpolar, or ionic. • P-Cl polar 2. Cl-Cl nonpolar 3. Ca-F ionic

22. Polar Bonds Partial positive charge Partial negative charge F has a higher electronegativity than H. H and F are sharing the electrons, but F has a greater pull on the shared electrons. F wins the “tug of war” and pulls the electrons closer giving F a partial negative charge. The electrons are pulled away from the H atom. This gives H a partial positive charge. These are NOT complete charges like in an ionic bond.

23. Polar, Nonpolar and Ionic Bonds

24. Video https://www.youtube.com/watch?v=_M9khs87xQ8

25. PROPERTIES OF COVALENT COMPOUNDS • Solids, liquids and gases • Lower melting points than ionic compounds • Nonconductors (nonelectrolytes) • All nonmetals • Ex: C12H22O11 (solid) H2O (liquid) CO2 (gas)

26. Octet Rule • Octet rule- atoms form bonds to have 8 valence electrons • Exception: H (2 e- ) • B (6 e-)

27. Lewis Dot Structures • F2

28. MATH METHOD • Count the total # of e- needed to satisfy the octet rule. (NEED) • Count the total # of valence e-. (HAVE) • Subtract #2 from #1 and divide by 2. This equals the # of bonds. • Make sure every atom obeys the octet rule.

29. Bond Energy • BOND ENERGY – energy required to break a bond. • Triple bond has the _____________ bond energy and a single bond has the ____________ bond energy. • Triple bond has the _____________ bond length and a single bond has the ____________ bond length.

30. Resonance • Resonance Structures- occur when more than 1 Lewis structure can be drawn for a molecule. • Ex: CO3-2 • SO2

31. Exceptions • Exceptions to the octet rule: • H ______ B_______ • Be __________ • Expanded Octet: _______than 8 e- on the central atom (Honors)

32. Examples (Honors) Ex: PCl5 SF6

33. Video http://ed.ted.com/lessons/what-is-the-shape-of-a-molecule-george-zaidan-and-charles-morton

34. MOLECULAR SHAPES • VSEPR MODEL • Bonds are made up of e- • Bonds repel each other • Bonds will spread outaround the central atom to be as far apart as possible • Lone pairs of electrons repel more than bonds repel each other • Count the number of bonded atoms and the number of lone pairs on the CENTRAL atom to determine the shape.

35. Molecular Shapes • Two atoms bonded to central atom. Can be single, double, or triple bonds. Ex: BeH2 Shape: linear Bond angle: 180°

36. Linear

37. Linear

38. Molecular Shapes 2. Three atoms bonded to central atom. Ex: BH3 Shape: trigonal planar Bond angle : 120°

39. Trigonal Planar

40. Trigonal Planar Ex: CH2O

41. Molecular Shapes 3. Two atoms and one or two lone pairs on the central atom. Ex. H2O Shape: bent Bond angle: 104.5°

42. Bent

43. Bent Ex: SO2

44. Molecular Shapes 4. Three atoms and one lone pair on the central atom. Ex: NH3 Shape: Trigonal Pyramid Bond angle: 107°

45. Trigonal Pyramid

46. Molecular Shapes 5. Four atoms around the central atom. Ex: CH4 Shape: Tetrahedral Bond angle: 109.5°

47. Tetrahedral

48. (Honors) Additional Shapes 6. Five atoms around the central atom. Ex: PCl5 Shape: TrigonalBipyramidal Bond angle: 90° and 120°

49. TrigonalBipyramidal

50. One Last Shape 7. Six atoms around the central atom. Ex: SF6 Shape: Octahedral Bond angle: 90°