1. Chapters 4 & 5 Chemical Bonding

2. Valence Electrons Outermost electrons s and p electrons for main group elements Responsible for chemical properties of atoms Participate in chemical reactions Core Electrons Valence Electron

4. Problems Write out the electron configurations for the following elements and identify how many core and valence electrons each has. 1)Mg 2)S 3)Br 4)Kr

5. Lewis Dot Structures LDS: a representation of an atom using its chemical symbol surrounded by dots that signify valence electrons

6. Problems Write the Lewis Dot Structures for the following atoms Li Be Br C N Ne

7. Li: [He]1s 1 Na: [Ne]2s 1 K: [Ar]3s 1

8. Octet Rule Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells –Natural electron configuration of the Noble Gases –Done by gaining, losing, or sharing electrons –Increases stability –H and He seek a “ Duet ”

9. Ionic Bonding Ions: atoms that have a charge due to gain or loss of electrons –Anion: (-) charged atom –Cation: (+) charged atom Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

10. Formula Unit

13. Ionic Compounds: compounds composed of oppositely charged ions that are held together by their attraction to each other Metal + Non-metal –NaCl Metal + Polyatomic Ion –NaNO 3 Polyatomic Ion + Non-metal –NH 4 Cl Polyatomic Ion + Polyatomic Ion –NH 4 NO 3 Net charge on compound equal to zero

15. Oxyanions SO 4 2- Sulfate SO 3 2- Sulfite PO 4 3- Phosphate PO 3 3- Phosphite NO 3 - Nitrate NO 2 - Nitrite ClO 4 - Perchlorate ClO 3 - Chlorate ClO 2 - Chlorite ClO - Hypochlorite

16. Rules For Naming Ionic Compounds 1)Name the cation by its elemental/polyatomic name 2)If the metal is a transition metal with a variable charge, indicate its charge with a Roman Numeral in parentheses 3)Next, name the anion and change its ending to “ -ide ” 4)If the anion is polyatomic, do not change the ending to “ -ide ” 5)Do NOT use prefixes (mono, di, tri etc.) to indicate how many of each atom are present

17. Problems Write the name for the following compounds: 1)KI 2)MgBr 2 3)Al 2 O 3 4)FeCl 2 5)CaSO 4 6)Ba(NO 2 ) 2 7)Cu(NO 3 ) 2

18. Write the Formula for the following ionic compounds: 8)Sodium Fluoride 9)Calcium Sulfite 10)Calcium Chloride 11)Iron (III) Oxide 12)Cobalt (II) Hydroxide 13)Ammonium Bromide 14)Ammonium Carbonate 15)Aluminum Carbonate

19. Iron (II) Chloride Iron (III) Chloride

20. Covalent Compounds Covalent Compounds: compounds composed of atoms bonded to each other through the sharing of electrons Electrons NOT transferred No + or – charges on atoms Non-metal + Non-metal Also called “ molecules ” Examples : –H 2 O –CO 2 –Cl 2 –CH 4

21. or H-H or Duet

22. Naming Covalent Compounds 1)Name the first non-metal by its elemental name 2)Add a prefix to indicate how many 3)Name the 2 nd non-metal and change its ending to “ -ide ” 4)Add a prefix to indicate how many 16 27 38 49 510

23. Problems Write the name of the following compounds: 1)CO 2)NI 3 3)N 2 O 4)SF 6 5)B 2 O 3

24. Write the formula for the following compounds: 6)Phosphorous Pentachloride 7)Nitrogen Monoxide 8)Dinitrogen Tetroxide 9)Tetraphosphorous Decoxide

25. Problems 1)KCl 2)Na 2 S 3)H 2 O 4)SO 2 5)K 3 PO 4 6)FeCl 3 7)(NH 4 ) 2 SO 4 8)SCl 2 9)Cu(OH) 2 10)P 2 O 5

26. 8)Sodium Iodide 9)Aluminum Sulfate 10)Phosphorous Pentabromide 11)Magnesium Nitride

27. Naming Acids Acids that do not contain oxygen 1)Begin the name with “ hydro ” 2)Name the anion, but change the ending to “ -ic ” 3)Add “ acid ” on the end HCl HF

28. Acids that contain oxygen 1)Do not put “ hydro ” at the beginning 2)Begin the name with the anion 3)If the anion has the ending “ -ate, ” change this to “ -ic acid ” 4)If the anion has the ending “ -ite, ” change this to “ -ous acid ” HClO 4 HClO 3 HClO 2 HClO

29. Problems Name the following 1)HBr(g) 2)HBr(aq) 3)HNO 2 (aq) 4)HNO 3 (aq) 5)HI (aq) 6)HI (g) 7)H 2 CO 3 (aq) 8)H 3 PO 4 (aq) 9)H 3 PO 3 (aq) 10)HCN (aq)

30. Molecular Structures

31. Water Ball & Stick ModelsSpace-Filling Models Methane

32. Ethanol

34. Lewis Dot Structures 1)Count the total number of valence electrons in the molecule. Ex: PCl 3 2)Use atomic symbols to draw a proposed structure with shared pairs of electrons. Atoms don’t tend to bond to other atoms of the same element when they can avoid it Exception: Carbon

35. 3)Place lone pair electrons around each (except H) to satisfy the octet rule, beginning with the terminal atoms 4)Place any leftover electrons on the central atom 5)If the number of electrons around the central atom is less than 8, change single bonds to the central atom to multiple bonds (double or triple). Ex: CH 2 O

36. Problems Draw the LDS’s for the following molecules: 1)Cl 2 O 2)C 2 H 4 3)C 2 H 6 O

37. What Things Like To Do Halogens –Like to be terminal –Like to have one single bond and 3 lone pairs (non-bonding electrons) Carbon –Likes to have 4 single bonds and no lone pairs A double bond counts as two singles A triple bond counts as three singles –Likes to be central –Likes to bond to other carbons

38. Silicon –Likes to do what carbon does Oxygen –Likes to have two single bonds and 2 lone pairs Sulfur –Likes to do what oxygen does Nitrogen –Likes to have 3 single bonds and one lone pair

39. Phosphorous –Likes to do what nitrogen does Hydrogen –Likes to be terminal with only one single bond –No lone pairs!

40. Problems 1)SH 2 2)C 3 H 8 3)Si 2 H 6 4)PI 3 5)CH 3 OH 6)C 2 H 2

41. 7)CCl 2 O 8)N 2 H 4 9)CH 2 OS 10)C 2 H 6 O 11)CO 12)BrHO

42. Electronegativity The measure of the ability of an atom to attract electrons to itself –Increases across period (left to right) and –Decreases down group (top to bottom) –fluorine is the most electronegative element –francium is the least electronegative element

44. Electronegativity Scale

45. Types of Bonding 1)Non-Polar Covalent Bond: Difference in electronegativity values of atoms is 0.0 – 0.4 Electrons in molecule are equally shared Examples: Cl 2, H 2, CH 4 EN Cl = 3.0 3.0 - 3.0 = 0 Pure Covalent

46. 2)Polar Covalent Bond: Difference in electronegativity values of atoms is 0.4 – 2.0 Electrons in the molecule are not equally shared The atom with the higher EN value pulls the electron cloud towards itself Partial charges Examples: HCl, ClF, NO EN Cl = 3.0 EN H = 2.1 3.0 – 2.1 = 0.9 Polar Covalent

47. 3)Ionic Bond: Difference in EN above 2.0 Complete transfer of electron(s) Whole charges EN Cl = 3.0 EN Na = 1.0 3.0 – 0.9 = 2.1 Ionic

49. Problems Predict the type of bonding in the following compounds using differences in EN values of the atoms. Indicate the direction of the dipole moment if applicable 1)KBr 2)HF 3)BrI 4)FI

50. Valence Shell Electron Pair Repulsion Theory VSEPR theory: –Electrons repel each other –Electrons arrange in a molecule themselves so as to be as far apart as possible Minimize repulsion Determines molecular geometry

56. Defining Molecular Shape Electron pair geometry: the geometrical arrangement of electron groups around a central atom –Look at all bonding and non-bonding e - ’ s Molecular Geometry: the geometrical arrangement of atoms around a central atom –Ignore lone pair electrons

59. 2 e - groups surrounding the central atom –e- pair geometry: linear –MG: linear –AXE designation: AX 2 E 0 A: Central Atom X: Bonding pairs E: Non-bonding pairs –Example: BeCl 2

60. 3 e - groups 3 Bonds, 0 Lone Pairs –e - PG: Trigonal Planar (Triangular planar) –MG: Trigonal Planar –AX 3 E 0 –BF 3 2 Bonds, 1 Lone Pair – e - PG: Trigonal Planar (Triangular planar) –MG: Bent/angular –AX 2 E 1 –GeCl 2

61. 4 e - groups 4 bonds, 0 Lone Pairs –e - PG: Tetrahedral –MG: Tetrahedral –AX 4 E 0 –CH 4 3 bonds, 1 Lone Pair –e - PG: Tetrahedral –MG: Triangular Pyramidal –AX 3 E 1 –NH 3

62. 2 bonds, 2 Lone Pairs –e - PG: Tetrahedral –MG: Bent/Angular –AX 2 E 2 –H 2 O

63. Drawing LDS With Correct Geometry

64. Molecular Polarity

66. Problems 1)NF 3 2)CH 2 O 3)CBr 4 4)CHCl 3 5)CH 2 Cl 2 Draw the Lewis Dot Structures for the following molecules and then identify the direction of polarity, if any.