1. Chapter Four: Forces Between Particles 2, 12, 14, 20, 22, 26- 32, 36, 38, 48-58, 62, 66-74

2. Chemical Bonding Review Compounds and Molecules are held together by chemical bonds Three types of bonds –Ionic Metals and non-metals –Covalent Non-metal and Non-metal –Metallic Between atoms of metals

3. Octet Rule All atoms strive to have electronic configurations like the Noble Gases Eight electrons in the outermost shell, highest principle quantum number (n) Except H and He follow duet rule –Want two electrons in outermost shell How do the atoms achieve an octet?

4. Taking or Giving and Sharing Electrons Ionic Bonds –Atoms take or give electrons from other atoms Covalent Bonds –Atoms share electrons between themselves Metallic Bonds –Sea of electrons

5. Valence Electron Review 2

6. Lewis Dot Structures For Atoms/Ions Symbol represents the nucleus and all electrons except for those in the valence shell Give the Lewis Dot Structure for: NaFO 2- Species with the same number of electrons are isoelectric O 2- F - NeNa + Mg 2+ How many electrons does each species have?

7. Lewis Dot Structures GN Lewis developed the theory of covalent bonding Structures showing covalent bonds are called Lewis structures Each line represents a shared pair of electrons (2 electrons) Lone pairs of electrons are shown by a pair of dots

8. Drawing Lewis Structures Decide on atom connectivity and placement –Hydrogen (never in the middle) is frequently bonded to oxygen –Oxygen is rarely the central atom –Oxygen will not bond to oxygen (except O 2 or O 3 ) –Carbon will be the central atom –Least electronegative atom is in the middle

9. Drawing Lewis Structures Count the total number of valence electrons –An atom’s number of valence electrons is equal to its group number Determine the total number of shared electrons electrons needed – valence electrons present Connect the atoms with single bonds –A single bond is one shared pair of electrons Use lone pairs and/or multiple bonds to give each atom an octet of electrons

10. Lewis Structure (Single Bonds) Draw Lewis Structures for: H 2 O HCl NH 3

11. Lewis Structures (Multiple Bonds) CO 2 N 2

12. Ions Definition: Ions are atoms or groups of atoms with an electrical charge Cations: are positively charged, due to loss of electrons (Metals) Anions: are negatively charged, due to gain of electrons (Non-Metals) Number of electron’s gained or loss is due to atoms wanting Octet

13. Examples of Ions Na Ra Al Se O Cl F

14. Ionic Compounds Ionic compounds are held together by ionic bonds, or the attraction of oppositely charged ions In the solid state, ionic compounds form crystalline lattices –Cations are attracted to all the neighboring anions, not just one –Thus, there are no discrete ionic “molecules” Ball and Stick Model

15. Transition Metal Cations Most transition metals form more than 1 cation +1 onlyAg + +2 onlyZn 2+, Cd 2+ +1 and +2Hg 2 2+, Cu + Hg 2+, Cu 2+ +2 and +3Cr 2+, Fe 2+, Co 2+ Cr 3+, Fe 3+, Co 3+ +2 and +4Sn 2+, Pb 2+ Sn 4+, Pb 4+

16. Polyatomic Ions NO 3 - nitrate SO 4 2- sulfate PO 4 3- phosphate NO 2 - nitrite SO 3 2- sulfite HPO 4 2- monohydrogenphosphate CO 3 2- carbonate NH 4 + ammonium H 2 PO 4 - dihydrogenphosphate HCO 3 - Bicarbonate Or hydrogen carbonate OH - hydroxide C 2 H 3 O 2 - acetate

17. Formulas of Ionic Compounds The net charge on a formula unit must be zero  (+) charges =  (-) charges Since there are no ionic “molecules” the formula of an ionic compound is the simplest ratio of cation to anion that gives an electrically neutral combination Al 3+ and O 2- Ca 2+ and O 2-

18. Writing Ionic Compound Formulas Write the formula for each of the following pairs of ions Na and Oxygen Mg and Fluorine Rb and Iodine

19. Nomenclature Rules for naming compounds and molecules Anions –Name the element, drop the ending leaving the root and add “ide” Element – root + “ide” Cl O N S I

20. Naming Ionic Compounds 1.Name the cation by naming the element If the cation is a transition metal you need to distinguish the charge using Roman Numerals Fe 2+ is named Iron (II) Pb 4+ is named Lead (IV) 2.Name the anion Can be an elemental anion or polyatomic 3.Combine them as two words

21. Naming Ionic Compounds K 2 O Li 2 CO 3 K 2 SO 4 NaHCO 3 Cr 2 O 3

22. Formulas from Names What are the formulas of these compounds? calcium sulfide iron (III) acetate Chromium (III) sulfate

23. Naming Molecular Compounds Name each element Indicate how many of each element is present with a prefix multiplier –Mono =1; di =2; tr i=3; tetra =4; penta =5; hexa =6; hepta = 7; octa = 8; nona = 9; deca = 10 Add the suffix “ide” to the last element The prefix multiplier mono is left off of the first element in the compound

24. Naming Molecular Compounds: Examples IBr NI 3 N 2 O 4

25. Formulas from Names Sulfur dioxide Diphosphorous pentoxide Carbon tetrachloride

26. Molecular Compounds: Common Names These compounds have common (non-systematic names) –Water (H 2 O) Dihydrogen monoxide –Ammonia (NH 3 ) Nitrogen trihydride –Methane (CH 4 ) Carbon tetrahydride –Nitrous oxide (N 2 O) Dinitrogen monoxide –Hydrazine (N 2 H 4 ) Dinitrogen tetrahydride

27. Acids Acids are compounds that can donate an hydrogen ion (H + ion) Acids fall into two categories –Binary Acids HX –Oxoacids HXO n Polyatomic anions

28. Binary Acids Most binary acids result from dissolving the corresponding molecular compound in water Binary acids are named as hydro (stem name of X) ic acid HCl (g) HCl (aq) HCN (g) HCN (aq)

29. Oxoacids Oxoacids are named based on the oxoanion “Ate” anion => ic acid “Ite” anion => ous acid

30. Oxoacids CO 3 2- (carbonate anion)H 2 CO 3 (carbonic acid) NO 2 - (nitrite anion)HNO 2 (nitrous acid) NO 3 - (nitrate anion)HNO 3 (nitric acid) PO 4 3- (phosphate anion)H 3 PO 4 (phosphoric acid) SO 3 2- (sulfite anion)H 2 SO 3 (sulfurous acid) SO 4 2- (sulfate anion)H 2 SO 4 (sulfuric acid) Polyatomic Anion

31. Review of What We Know We can write formulas We can name compounds and molecules We can draw Lewis Structures –But what do these molecules look like?

32. VSEPR Theory VSEPR: Valence Shell Electron Pair Repulsion Like charges repel and want to be as far apart as possible Therefore a given combination of electrons will form into a specific shape

33. VSEPR 1.Draw the Lewis Structure 2.Assign the central atom (A) 3.Determine the number (n) of atoms bonded to (A) designate them (X n ) 4.Determine the number of lone pairs on (A) designate them (E m ) 5.Put together the AX n E m notation

34. X + E = 2 X + E = 3 X + E = 4 X + E = 5 X + E = 6

35. VSEPR Examples What is the geometry of CO 2 BF 3 H 2 O NH 3

36. Electronegativity Linus Pauling developed the electronegativity scale Electronegativity is a measure of an atom’s affinity for electrons Fluorine is the most electronegative element (EN=4.0) The closer an atom is to fluorine, the more electronegative it is

38. Polar Covalent Bonds If two atoms of identical electronegativity are bonded together, the bond is non-polar If two atoms of different electronegativity are bonded together, the bond is polar, and the electrons spend more time around the more electronegative atom –This creates partial charges The greater the difference in EN between two atoms, the more polar the bond –The limiting example of this is the ionic bond

39.  EN Type of Bonding 0.0Pure covalent bond (equal sharing of e - ‘s) 0.1 – 0.4Non-polar covalent bond (almost equal attraction for shared e - pairs) 0.5 – 1.4Polar covalent bond (unequal sharing of e - ’s) 1.5 – 3.2Ionic bond (e - transfer)

40. Example The bond in hydrogen is The bond in hydrogen chloride is

41. Molecular Polarity Bond dipoles are vectors The vectoral sum of the bond dipoles gives the molecular dipole Based on the shape of the molecules you can predict if the dipoles will cancel each other or if they will create a dipole moment If a dipole moment exists then the molecule is said to be polar If no dipole moment exists then the molecule is said to be non-polar

42. Molecular Polarity Examples Is carbon dioxide polar or non-polar? Is water polar or non-polar? Is boron trifluoride polar or non-polar?

43. Intermolecular Forces These are attractive forces between molecules or atoms or ions Immensely important –These forces hold DNA molecules in a helix and and are the mechanism for DNA transcription

44. Dipole Dipole Attraction This is the attraction between the opposite (partial) charges of polar molecules HCl  HCl  

45. Hydrogen Bonding This is generally stronger than dipolar attractions Hydrogen bonding occurs between a hydrogen atom and O, N or F. For H-bonding to happen the H must be directly bonded to a O, N or F. O—H H H This is an attraction not really a bond

46. Also called Van der Waal’s forces, these are created by instantaneous dipoles London forces are much weaker than either dipole-dipole or H-bonding London forces get stronger with larger atoms/molecules London Forces

47. He London Forces Between Helium Atoms He  He   He   He At a given instant, the electrons on an atom may be non-symmetrically distributed. This leads to creation of a temporary dipole. As the electrons re-distribute, the dipoles and and the attraction vanishes. This dipole induces temporary dipoles on neighboring atoms. For the merest fraction of time, there is a dipole-dipole attraction between the atoms.

48. Ion Dipole Attraction This is the attraction between an ionic charge and a polar molecule This attraction allows ionic solids to dissolve in water The strength of this force varies widely and depends on the magnitude of the dipole moment of the polar species and the size of the ion

49. A Sodium Ion and a Chloride Ion Hydrated by Water Molecules

50. Effects of Intermolecular Forces More intermolecular forces mean: –Higher boiling and melting points –More viscous liquids IM Forces also affect solubility –‘like dissolves like’

51. Predicting Boiling Points based on IMF’s SnH 4, CH 4, GeH 4, SiH 4 HBr, HI, HCl, HF

52. Trends in Boiling Point H2OH2O H2SH2S H 2 Se H 2 Te

53. Example Is carbon dioxide soluble in water? Explain

54. Example Are ionic compounds more soluble in water or in gasoline (a non-polar solvent)? Explain