1. Chemical Bonding

2. What is a Bond? Force that holds atoms together Results from the simultaneous attraction of electrons (-) to the nucleus (+)

3. Breaking/Forming Bonds When a bond is broken energy is absorbed –Endothermic When a bond is formed energy is released –Exothermic The greater the energy released during the formation of the bond, the greater its stability –Stable bonds require a great deal of energy to break

4. Lewis Dot Diagrams Use dots to represent the number of valence electrons How to write: –Write the symbol. –Put one dot for each valence electron –Electrons go on the 4 sides, no more than 2 per side

5. Dot Diagram Examples: Draw dot-diagrams for the following 1.Mg 2.C 3.Ne

6. Dot Diagrams - Ions For ions, use brackets and place the charge outside the brackets Examples: 1.Na + 2.O 2- 3.H +

7. Octet Rule Atoms will gain or lose electrons in order to have a full valence shell – like the nobles gases “Take the shortest route” Metals lose electrons to form positive ions (Cations) Nonmetals gain electrons to form negative ions (Anions)

8. Exceptions 1 st principle energy level only holds 2 electrons Transition elements can lose valence (s) and inner (d) electrons – this is why they have multiple oxidation states Some atoms may be stable with less than an octet – many compounds with B Some atoms may be stable with more than an octet – elements beyond period 2, especially P and S, the additional electrons are added to the d sublevel Molecules with an odd number of electrons – they will be unstable

9. Types of Bonds Ionic - Electrons are transferred from a metal to a nonmetal Covalent - Electrons are shared between 2 nonmetals –Polar Covalent – electrons are shared unequally –Nonpolar Covalent – electrons are shared equally Metallic - Electrons are mobile within a metal, “Sea of Electrons”

10. Identifying Bond Type Ionic – metal and a nonmetal Covalent – 2 nonmetals –Nonpolar Covalent (equal) Same atoms (diatomics, triatomics) –Polar Covalent (unequal) Unequal electronegativity Metallic – metals

11. Identifying Bond Types Indicate the type of bond present in each: 1.HCl 2.CCl 4 3.MgCl 2 4.O 2 5.Hg 6.H 2 O

12. Ionic Bonds Transfer of 1 or more electrons from a metal to a nonmetal The greater the electronegativity difference between atoms, the greater the ionic character Example: Sodium Chloride (NaCl) Na Cl X Na electron transferred to Cl

13. Monatomic Ions One atom in an ion Look at the valence electrons to determine the charges Examples: K +, O 2-

14. Polyatomic Ions More than one atom in the ion Reference Table E Charge belongs to the entire ion, not an individual atom Within the polyatomic ion the atoms are held together by covalent bonds When writing it, place ( ) around the entire ion, with the charge outside Examples: (NH 4 ) +, (H 3 O) +, (CO 3 ) 2-

15. Writing Ionic Formulas You need an equal amount of positive and negative charges, so that the compound is neutral Ionic Formulas are always written as empirical formulas (reduced)

16. Examples 1.Na 1+ + Cl 1- 2.Mg 2+ + Cl 1- 3.Ca 2+ + CO 3 2- 4.Al 3+ + O 2-

17. Criss Cross Method 1.Write the symbol for the cation and anion 2.Write each ion’s charge as a superscript 3.Criss-cross the charges to become subscripts of the other ion  Do not put (+) or (-) charges in the final formula 4.Reduce to least common multiple (empirical formula)

18. Ionic Formulas Write the formula for the compound formed from the following ions: 1.Mg 2+ + Cl - 2.Ca 2+ + CO 3 2- 3.Al 3+ + O 2- 4.Calcium ion + hydroxide ion

19. Naming Ionic Compounds Name the cation first, the anion second Cation keeps its name, anion changes its ending to –ide (Chlorine → Chloride) Do not change the ending of polyatomic ions Examples: 1.NaCl 2.CaCO 3 3.MgF 2

20. Stock System – only used for positive ions Some cations have more than one positive oxidation states A roman numeral is used to indicate the charge of the positive ion

21. Stock System Examples 1.Iron (II) Chloride 2.Iron (III) Oxide 3.Copper (II) Oxide 4.a. What charge does copper have in copper II sulfate? b. What is the formula for copper II sulfate?

22. Ionic Salts Salts are ionic compounds made up of cations and anions The ratio of cations to anions is always such that an ionic compound has no overall charge Many of the ions are bonded together to form a crystal

23. Properties of Ionic Salts Ionic Bonds are very strong Very high melting and boiling points –All solids at STP Hard Brittle

24. Melting and Boiling Points of Compounds Compound Name FormulaType of Compoundmp ( o C) bp ( o C) Magnesium FlourideMgF 2 Ionic12612512 Sodium ChlorideNaClIonic8011686 Calcium IodideCaI 2 Ionic7841373 Iodine MonoChlorideIClCovalent27370 Carbon tetrachlorideCCl 4 Covalent-23350 Hydrogen FlourideHFCovalent-83293 Hydrogen SulfideH2SH2SCovalent-86212 MethaneCH 4 Covalent-182109

25. Properties of Salts (cont’d) Do not conduct electricity as solids Do conduct electricity when the salt melts or is dissolved in water (liquid phase or aqueous) –In order to conduct electricity a substance must have free moving charged particles –In the solid phase the ions are not free to move

26. Covalent Bonds Sharing of electrons between 2 nonmetals

27. Non-Polar Covalent Electrons are shared equally Equal distribution of electrons All diatomic molecules have non-polar covalent bonds

28. Nonpolar Covalent Examples 1.Flourine (F 2 ) 2.Hydrogen (H 2 )

29. Polar Covalent Unequal Sharing of electrons Unequal distribution of electrons –Partial positive and partial negative charges –The side with the higher electronegativity will have a greater share of the electron(s) resulting in a partial negative charge The greater the electronegativity difference, the more polar the bond is

30. Polar Covalent Examples 1.HCl 2.H 2 O

31. Dipoles Form when the charge in a bond is asymmetrical –Present in polar bonds –Partial positive and partial negative charges

32. Polar Bonds / Dipoles Isn’t a whole charge just a partial charge means a partially positive means a partially negative Example: H - Cl +---→ The Cl pulls harder on the electrons (more eneg) The electrons spend more time near the Cl   

33. Dipole Examples 1.Which molecule contains more polar bonds? a. CCl 4 b. CH 4 2. Which has a stronger dipole? a.HCl b.HBr

34. Properties of Molecular Substances (Covalent Compounds) Soft Low melting points and boiling points –Many exist as gases or liquids at STP Poor conductors of heat and electricity (in all phases) Examples: H 2 O, CCl 4, NH 3, C 6 H 12 O 6, O 2

35. Molecular Formulas (Covalent Compounds) Contain covalent bonds Tells you how many atoms are present in a single molecule Named similarly to ionic compounds, except use prefixes to indicate the number of atoms per molecule

36. Prefixes Mono- is only used for the second element –Example: CO = carbon monoxide Mono-1Hexa-6 Di-2Hepta-7 Tri-3Octa-8 Tetra-4Nona-9 Penta-5Deca-10

37. Examples 1.CCl 4 2.H 2 O 3.NO 4.N 2 O 5 5.BBr 3

38. Structural Formulas Specifies how atoms are bonded together Dashes represent bonds 2 atoms can share up to 3 pairs of electrons

39. Single Bonds 2 atoms share 1 pair of electrons (2 electrons) Examples: 1.Ammonia (NH 3 ) 2.Chlorine (Cl 2 ) 3.Hydrochloric Acid (HCl)

40. Double Covalent Bonds 2 atoms share 2 pairs of electrons (4 electrons) 2 bonds between 2 atoms Examples: 1.Carbon Dioxide (CO 2 ) 2.Oxygen (O 2 )

41. Triple Covalent Bond 2 atoms share 3 pairs of electrons (6 electrons) 3 bonds between 2 atoms Examples: 1.Nitrogen (N 2 ) 2.Ethyne (C 2 H 2 )

42. Bond Length/Strength Length: –Single > Double > Triple The more electrons in a bond, the greater the attraction, therefore shorter –As you move down a group bond length increases Due to increasing molecular size Strength: –Triple is the strongest, most stable, requires the most energy to break

43. Network Solids Covalently bonded atoms are linked into a giant network (macromolecules) Examples: Diamond (C), Graphite (C), Silicon Carbide (SiC), and Silicon Dioxide (SiO 2 )

44. Network Solids Properties: –Hard –High melting and boiling points –Do not conduct heat and electricity

45. Metallic Bonding Sea of Electrons –Electrons are free to move through the solid. ++++ ++++ ++++

46. Properties of Metallic Solids Very Strong Good conductors of heat and electricity because electrons are free to move about Luster High melting point (except Hg) –All solids at STP (except Hg) Malleable, Ductile

47. VSEPR Theory In a small molecule, the electron pairs are as far away from each other as possible –VSEPR = Valence Shell Electron Pair Repulsion

48. Linear Drawn on a straight line All molecules of only 2 atoms are linear Many 3 atom molecules are linear, if there are no unshared electron pairs on the central atom If both ends are the same, the molecule is nonpolar (Symmetrical = Nonpolar) If the ends are different, the molecule will be polar (Asymmetrical = Polar) Bond Angle = 180 o See Molecules Examples: H 2, CO 2, HCl

49. Tetrahedral

50. A central atom bonded to 4 other atoms 3-D shape allows the electron pairs to get as far away from each other as possible C HH H H 109.5º

51. Tetrahedral If all the ends are the same, NONPOLAR If the ends are different, POLAR Bond Angle = 109.5 o See Molecules Examples: 1. CH 4 2. CH 3 Cl

52. Pyramidial A central atom is bonded to 3 other atoms and the central atom has an unshared electron pair 3-D, like a pyramid Always POLAR Bond Angle = 107 o See Molecules Example: NH 3

53. Bent

54. A central atom is bonded to 2 other atoms and the central atom has 2 unshared electron pairs Always POLAR Bond angle = 105 o See Molecules Example: H 2 O

55. Intermolecular Attractions/Forces Forces between molecules Determines boiling point, melting point, vapor pressure, surface tension –The stronger the intermolecular attractions, the higher the boiling point All intermolecular attractions are weaker than actual bonds Polar molecules will have stronger IMFs than nonpolar molecules –The greater the polarity the stronger the IMF

56. Dipole-Dipole Forces Occurs between 2 polar molecules The positive end of one molecule is attracted to the negative end of another molecule The greater the electronegativity difference is, the more polar the bond will be and the stronger the dipole will be Example: HCl

57. Dipole Examples 1. Which would have the strongest intermolecular forces? Explain Why. a. HCl b. HBr 2. Which would have the weakest intermolecular forces? Explain Why. a. H 2 S b. H 2 O

58. Hydrogen Bonds Special, Strong type of Dipole Attractions Attraction of a covalently bonded H atom to a F, O, or N atom on another covalent compound H H O ++ -- ++ H H O ++ -- ++

59. Hydrogen Bonds VERY STRONG Molecules with H bonds will have high boiling points, melting points, and surface tension –See Water HeatingSee Water Heating Example: NH 3

60. H-bonds Examples 1. Which sample has Hydrogen Bonds? a. H 2 b. HF c. F 2 d. HCl 2. Rank in order from strongest (1) to weakest (3). a. Hydrogen Bonds b. Covalent Bonds c. Dipole-Dipole Attractions

61. Molecule – Ion Attractions Attraction between a polar compound and an ion (ionic salt) Polar substances (such as water) attract ions from ionic compounds in solution This allows ionic substances to dissolve in polar solvents (water) –The anion is attracted to the positive end of the polar solvent –The cation is attracted to the negative end of the polar solvent –The ion dissociates (falls apart) Example: NaCl(aq)

62. Molecule-Ion Examples 1. Molecule-Ion attractions are present in which sample? a. HCl(l)c. KCl(l) b. HCl(aq)d. KCl(aq) 2. When sodium chloride dissolves in water the chloride ion is attracted to a. The positive part of the water, the O atom b. The negative part of the water, the O atom c. The positive part of the water, the H atom d. The negative part of the water, the H atom

63. Van Deer Waals Forces Very weak Exist between non-polar molecules Caused by momentary dipoles Increases as molecular mass increases

64. VDW Examples 1. Rank in order from weakest to strongest: –Hydrogen Bonds –Covalent Bonds –Van deer Waals Forces –Dipole-Dipole Attractions 2. Which would have the strongest intermolecular forces? a. H 2 b. Cl 2 c. F 2 d. Br 2