How do you think H and O in water are bonded together? Draw a picture to help your explanation.

1. How do you think H and O in water are bonded together? • Draw a picture to help your explanation.

2. Ch. 6 Bonding 6.1 Introduction to Chemical Bonding

3. Chemical Bonds • atoms rarely exist alone • when atoms are bonded together, they have less potential energy and are more stable • What is potential energy? • chemical bond – mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

4. Ionic Bonds • results from electrical attraction between large numbers of cations and anions • atoms donate or accept electrons from each other

5. Covalent Bonds • results from sharing of electron pairs between two atoms • the electrons shared belong to both atoms

6. Ionic vs. Covalent

7. Ionic vs. Covalent • bonding usually does not fall in one category or the other, but somewhere in between • type of bond depends on the elements differences in electronegativities 0.3

8. Practice • Determine whether each of the following bonds will be: Ionic or covalent

9. Patterns • What kind of patterns do you see? • metals + nonmetals = ionic • nonmetals + nonmetals = covalent

10. Ch. 6 Bonding 6.2 Covalent Bonding

11. What is the difference between covalent and ionic bond? How do you determine which a compound contains?

12. Molecular Compounds molecule: neutral group of atoms held together by covalent bonds molecular compound: compound whose simplest unit is a molecule

13. Formulas chemical formula: tells the number of each type of atom in a compound molecular formula: tells the number of each type of atom in a molecular compound ex. H2O, Cl2, C6H12O2

14. Molecular Compounds diatomic molecule: a molecules containing only 2 atoms usually refers to 2 of the same atoms ex: O2, Br2, F2, etc. 7+1 rule

15. Formation of Covalent Bond

16. Formation of Covalent Bond approaching nuclei and electron clouds are attracted to each other to create a decrease in PE two nuclei and two electron clouds repel each other creating an increase in PE

17. Formation of Covalent Bond a distance between the nuclei is reached in which repulsion and attraction forces are equal potential energy is at the lowest point possible at the bottom of the curve on PE graph

18. Covalent Bonds Bond Length distance between two bonded atoms at their lowest PE average distance since there are some vibrations measured in pm (1012 pm = 1 m) stronger the bond, shorter the bond

20. Covalent Bonds Bond Energy energy is released when atoms become because they have lower PE the same amount of energy must be used to break the bond and form neutral isolated atoms stronger bond, higher bond energy average since varies a small amount based on atoms in entire molecule in kJ/mol

22. Which elements naturally exist as diatomic molecules? Remember, the 7 + 1 rule How many valence electrons do each of the halogens have?

23. Octet Rule representative elements can “fill” their outer energy level by sharing electrons in covalent bonds Octet Rule- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons Duet Rule- applies to H and He

24. Octet Rule Less than 8: Boron: 6 in outer energy level More than 8: anything in 3rd period or heavier because may use the empty d orbital ex: S, P, I

25. Electron Dot Diagrams a way to show electron configuration identifies the number and pairing of valence electrons to show how bonding will occur write the noble gas notation identify the number of valence identify how many are paired and how many are alone do not go by Figure 6-10

26. Example Nitrogen 1s2 2s2 2p3 5 valence 2 are paired 3 are alone Sulfur 1s2 2s2 2p6 3s2 3p4 6 valence 4 paired (2 pairs) 2 are alone N

27. Lewis Structures like dot diagrams but for entire molecules atomic symbols represent nucleus and core electrons and dots or dashes represent valence electrons unshared electrons: (lone pairs) pair of electrons not involved in bonding written around only one symbol bonding electrons: written in between 2 atoms as a dash

28. Types of Bonds single- sharing of one pair of electrons weakest, longest double- sharing of 2 pairs of electrons stronger and shorter triple- sharing of 3 pairs of electrons strongest and shortest multiple bonds include double and triple bonds

29. Drawing Lewis Structures find the number of valence electrons in each atom and add them up draw the atoms next to each other in the way they will bond add one bonding pair between each connected atoms add the rest of the electrons until all have 8 (consider exceptions to octet rule)

30. Example 1 CH3Cl methyl chloride C: 4 x 1 = 4 H: 1 x 3 = 3 Cl: 7 x 1 = 7 total = 14 electrons carbon is central H H C Cl H duet octet duet octet duet H H C Cl H

31. Example 2 NH3 ammonia N: 5 x 1 = 5 H: 1 x 3 = 3 total = 8 N is central H N H H

32. Example 3 N2 nitrogen gas N: 5 x 2 = 10 10 electrons N N N N

33. Example 4 CH2O formaldehyde C: 4 x 1 = 4 H: 1 x 2 = 2 O: 1 x 6 = 6 total = 12 C is central H C H O

34. Example 5 O3 ozone O: 6 x 3 = 18 two completely equal arrangements the real structure is an average of these two where each bond is sharing 3 electrons instead of 4 or 2 O O O O O O

35. Resonance Structures resonance – bonding between atoms that cannot be represented in on Lewis structure show all possible structures with double-ended arrow in between to show that electrons are delocalized O O O O O O

36. Example 6 NO31- N: 5 x 1 = 5 O: 6 x 3 = 18 total = 23 + 1 = 24

37. Covalent Network Bonding a different type of covalent bonding not specific molecules lots of nonmetal atoms covalently bonded together in a network in all directions example: diamond silicon dioxide graphite

38. Ch. 6 Bonding 6.3 Ionic Bonding

39. Ionic Compounds ionic bonds do NOT form molecules chemical formulas for ionic compounds represent the simplest ratio of ion types made of anions and cations

40. Ionic Compounds combined so that amount of positive and negative charge is equal usually crystalline solid formula of ionic compound depends of the charges of the ions combined

41. Formation attractive forces: oppositely charged ions nuclei and electron clouds of adjacent ions repulsive forces: like-charged ions electrons of adjacent ions

42. Formation distance between the ions creates a balance between those forces ions minimize their PE by combining in an orderly arrangement called a crystal lattice

43. Formation specific lattice pattern created depends on: charges of ions size of ions Calcium Bromide: each Ca2+ is surrounded by 8 F- each F- is surrounded by 4 Ca2+ Sodium Chloride each Na+ is surrounded by 6 Cl- each Cl- is surrounded by 6 Na+

44. Ionic vs. Molecular ionic bonds and molecular bonds are both strong ionic bonds connect all ions together molecules are more easily pulled apart because intermolecular forces are weak

45. Ionic vs. Molecular Molecular Compounds: low melting and boiling points many are gases at room temperature Because the intermolecular forces of the molecules are weak so they are easily separated

46. Ionic vs. Molecular Ionic Compounds: higher melting and boiling points all are solid at room temperature hard:Because of the strong forces, it is difficult for one layer of ions to move past another brittle:if one layer is moved, the layers come apart completely

47. Ionic vs. Molecular Ionic Compounds: good conductors in liquid state Because ions are free to move and carry charge poor conductor in solid state Because ions are fixed in place

48. Polyatomic Ions charged group of covalently bonded atoms Example: CN-

49. NH4+ : ammonium ion

50. SO42- : sulfate ion 5 x 6 = 30 total = 30 + 2 = 32 OH- : hydroxide ion 6 + 1 + 1 = 8 total O O S O O O H