Covalent Bonding and Lewis Structures

1. Covalent Bonding and Lewis Structures

2. Day 1 • Lewis Dot Structures and Bond Polarity

3. Ionic Bonding • Generally occurs between metals and nonmetals • Can also occur with polyatomic ions.

4. Ionic Bonding • Involves the transfer of electrons, followed by electrostatic attraction.

5. Covalent Bonding • Generally occurs between nonmetals. • Involvessharingof electrons, rather than transfer.

6. Octet Rule • Atoms will acquire, through sharing or transfer, the electron configuration of a noble gas. This happens in orderfor the atoms to gain stability. • Most noble gases have 8 valence electrons • He is the exception

7. Octet Rule • These elements do not need a full octet: • H, He, Li, Be, B • Needs only 6 electrons to be stable Needs only two electrons to be stable

8. Valence Electrons • The electrons in the highest occupied energy level. There are two ways to determine the number of valence electrons. A. Using the electron configuration: 1s22s22p63s23p2, how many valence electrons are in this element? 2 + 2 = 4

9. Valence Electrons B. Look at the group number to determine valence electrons

10. Dot Models • The number of dots is equal to the number of valence electrons. 1 2 4 3 8 7 5 6 P Example: Phosphorus

11. Bonding in Covalent Molecules • Each atom in the molecule is connected by bonds. • Bonds are shared pairs of electrons, and they are represented by a dash. • Pairs that are not shared are called unshared electrons or lone pairs.

12. A single bond is created by one shared pair of electrons • A double bond is created by 2 shared pairs of electrons • A triple bond is created by 3 shared pairs of electrons.

13. Bond Polarity • In polarbonds, shared pairs of electrons are pulled between the nuclei of atoms sharing them. Sometimes electrons are pulled equally and sometimes they are not. This has to do with electronegativity. • Recall: electronegativity is the ability to attract electrons.

14. Nonpolar Covalent Bonds • The electrons are shared equally. • All diatomic elements are nonpolar.

15. Polar Covalent Bonds • The electrons are shared unequallydue to electronegativity • The more electronegative atom will have a stronger attraction for the bonded electrons and will have a slightly negative charge. • The less electronegative atom will have a slightly positive charge.

16. Example - HCl • Look up the electronegativity values for hydrogen and chlorine.

18. Example: HCl Which is more electronegative? • H: 2.1 • Cl: 3.0 Chlorine has a slightly negative charge, while hydrogen has a slightly positive charge.

19. There are two ways to communicate the polarity of HCl: H – Cl H – Cl  +  - The lowercase Greek letter delta shows that the atoms involved acquire only partial charges. The arrow points to the more electronegative atom.

20. Example: Water Is hydrogen or oxygen more electronegative?

21. Oxygen!!

22. Example: Water Is hydrogen or oxygen more electronegative? Oxygen!  - O H H The O-H bonds are polar.  +  +

23. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? • Br and Cl?

24. 3.0 – 2.1 = .9

25. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? • Br and Cl? Polar!

26. 4.0 – 3.0 = 1.0

27. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? • Br and Cl? Polar! Nonpolar!

28. 3.5 – 1.0 = 2.5

29. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? • Br and Cl? Polar! Nonpolar! Ionic!

30. 3.0 – 2.8 = .2

31. The difference in electronegativities indicates the type of bond the atoms will form. • What type of bond will form between: • N and H? • F and F? • Ca and O? Ionic! • Br and Cl? Nonpolar! Polar! Nonpolar!

32. Summary

33. Day 1 Homework! • Complete Page 5 • Extra Credit!!!!!!!!!!!!

34. Day 2: Covalent Bonding and Drawing the Lewis Structures of Molecules Sometimes a picture of a molecule is helpful in determining its structure looks like. For example is CH4 C-H-H-H-H or H-C-H-H-H or H-H-C-H …… How do we know? Methane looks like….

36. Steps for Drawing Lewis Dot Structures: • Look at your chemical formula and arrange your atoms as they most likely be positioned using the following rules: • The first atom in your formula will be your central atom unless it is hydrogen. Hydrogen is never central – use the next element in the formula if hydrogen is listed first. • Example: CH4 H H C H H

37. H 2. Add your valence electrons to each atom. 3. Draw a skeleton of the molecule, connecting atoms in the molecule with single bonds (1 shared pair) • Make sure each atom is happy even if they have to share. Each atom wants to have eight electrons around it, two to each side of the box. Each bond, or stick, counts as two electrons for each atom. C

38. Drawing CH4

39. Extra Practice: H2O HCP CO2 N2

40. 4. There are exceptions to the octet rule:

41. Day 2 Homework: page 9

42. Day 3 • Electron and Molecular Geometry • Once we have drawn the Lewis Structure, now we can describe how the atom looks in 3D.

43. Electron Geometry • The geometry based on the number of electron groups on the central atom. • It does not matter if the groups are a single bond, a multiple bond, or lone pairs.

44. Know the following Geometries

45. Molecular Shape • In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts whichever 3D geometry minimizesthis repulsion. • We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.

46. Know the following Molecular Geometries

48. Complete Day 3 Homework – page 12

49. Bond Polarity • In covalent bonds, shared pairs of electrons are pulled between the nuclei of atoms sharing them. • Sometimes electrons are pulled equally and sometimes they are not. • There are two types of covalent bonds: • Non-polar • Polar

50. Molecule Polarity • If a molecule has all nonpolarbonds then the molecule is nonpolar. • If a molecule has a polar bond then the whole molecule is usually polar, but not always.