3. Atom – the smallest unit of matter “indivisible” Helium atom

4. electron shells/orbits a)Atomic # = number of Protons & Electrons in a typical neutral atom b)Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells. a)1 st Orbit – 2 electrons3 rd Orbit – 18 electrons b)2 nd Orbit – 8 electrons4 th Orbit – 32 electrons c)Electron shells determine how an atom behaves (reacts) when it encounters other atoms

5. Valence Electrons: 1)Outer Orbit = Valence Electrons 1)Determine reactivity of elements 2)Can be gained, lost, or shared 3)Can’t have more than 8 electrons!

6. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons C would like to N would like to O would like to Gain 4 electrons Gain 3 electrons Gain 2 electrons

7. Why are electrons important? 1)Elements have different electron configurations 1)Exp. Lewis Dot Structures & Bohr Models 1)Na (Sodium) – Atomic # 11 or Na* 2)O (Oxygen) – Atomic #8 or ***O***  different electron configurations mean different levels of bonding

10. Electron Dot Structures (AKA Lewis Dot Structures) Symbols of atoms with dots to represent the valence-shell electrons 1 2 13 14 15 16 17 18 H  He:            Li  Be   B   C   N   O  : F  : Ne :                    Na  Mg   Al   Si   P   S  : Cl  : Ar :        

11. Learning Check A. X would be the electron dot formula for 1) Na2) K3) Al B. X would be the electron dot formula 1) B2) N3) P

12. How Electrons are used: 1.Bond: Force that holds groups of two or more atoms together and makes the atoms function as a unit. 1.Example: H-O-H 1.Bond Energy: Energy required to break a bond.

13. Chemical bonds: an attempt to fill electron shells 1.Ionic bonds 2.Covalent bonds 3.Metallic bonds

14. IONIC BOND bond formed between two ions by the transfer of electrons cation vs anion

15. Formation of Ions from Metals Ionic Compounds - A compound resulting from a positive ion (usually a metal) combining with a negative ion (usually a non-metal). : M + + X -  MX Ionic compounds result when metals react with nonmetals Metals lose electrons to match the number of valence electrons of their nearest noble gas Positive ions form when the number of electrons are less than the number of protons Group 1 metals  ion 1+ Group 2 metals  ion 2+ Group 13 metals  ion 3+

17. Formation of Sodium Ion Sodium atom Sodium ion Na  – e   Na + 2-8-1 2-8 ( = Ne) 11 p + 11 p + 11 e - 10 e - 0 1 +

18. Formation of Magnesium Ion Magnesium atom Magnesium ion  Mg  – 2e   Mg 2+ 2-8-2 2-8 (=Ne) 12 p + 12 p + 12 e- 10 e - 0 2 +

19. Some Typical Ions with Positive Charges (Cations) Group 1Group 2Group 13 H + Mg 2+ Al 3+ Li + Ca 2+ Na + Sr 2+ K + Ba 2+

20. Learning Check A. Number of valence electrons in aluminum 1) 1 e - 2) 2 e - 3) 3 e - B. Change in electrons for octet 1) lose 3e - 2) gain 3 e - 3) gain 5 e - C.Ionic charge of aluminum 1) 3- 2) 5- 3) 3 +

21. Solution A. Number of valence electrons in aluminum 3) 3 e - B. Change in electrons for octet 1) lose 3e - C.Ionic charge of aluminum 3) 3 +

22. Learning Check Give the ionic charge for each of the following: A. 12 p + and 10 e - 1) 02) 2+3) 2- B. 50p + and 46 e- 1) 2+2) 4+3) 4- C. 15 p + and 18e- 2) 3+ 2) 3-3) 5-

23. Ions from Nonmetal Atoms In ionic compounds, nonmetals in 15, 16, and 17 gain electrons from metals Nonmetal add electrons to achieve the octet arrangement Nonmetal ionic charge: 3-, 2-, or 1-

24. Fluoride Ion unpaired electronoctet     1 - : F  + e  : F :     2-7 2-8 (= Ne) 9 p+ 9 p + 9 e- 10 e- 0 1 - ionic charge

25. Ionic Bond Between atoms of metals and nonmetals with very different electronegativity Bond formed by transfer of electrons Produce charged ions at all states. Conductors and have high melting point. Examples; NaCl, CaCl 2, K 2 O

27. Ionic Bonds: One Big Greedy Thief Dog!

28. 1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.

29. COVALENT BOND bond formed by the sharing of electrons

30. Covalent Bond Between nonmetallic elements of similar electronegativity. Formed by sharing electron pairs Stable non-ionizing particles, they are not conductors at any state Examples; O 2, CO 2, C 2 H 6, H 2 O, SiC

32. Bonds in all the polyatomic ions and diatomics are all covalent bonds

33. Covalent Bonds Nonpolar –Electrons are shared equally –H 2, O 2 Polar –Electrons are not shared equally –H 2 O

34. when electrons are shared equally NONPOLAR COVALENT BONDS H 2 or Cl 2

35. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Molecule (O 2 ) Oxygen Molecule (O 2 )

36. when electrons are shared but shared unequally POLAR COVALENT BONDS H2OH2O

37. Polar Covalent Bonds: Unevenly matched, but willing to share.

38. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

39. Chemical bonds: an attempt to fill electron shells 1.Ionic bonds 2.Covalent bonds 3.Metallic bonds

40. METALLIC BOND bond found in metals; holds metal atoms together very strongly

41. Metallic Bond Formed between atoms of metallic elements Electron cloud around atoms Good conductors at all states, lustrous, very high melting points Examples; Na, Fe, Al, Au, Co

42. Metallic Bonds: Mellow dogs with plenty of bones to go around.

43. Ionic Bond, A Sea of Electrons

44. Metals Form Alloys Metals do not combine with metals. They form Alloys which is a solution of a metal in a metal. Examples are steel, brass, bronze and pewter.

45. Lewis Structures of molecules 1.Single Bond: Two atoms sharing one electron pair. Example: H 2 2.Double Bond: Two atoms sharing two pairs of electrons. Example: O 2 3.Triple Bond: Two atoms sharing three pairs of electrons. Example: N 2 Resonance Structures: More than one Lewis Structure can be drawn for a molecule. Example: O 3

46. Rules for Lewis structures of molecules 1.Write out valence electrons for each atom 2.Connect lone electrons because lone electrons are destabilizing 1.Become two shared electrons 1.Called a “bond” 3.Check to see if octet rule is satisfied 1.Recall electron configuration resembling noble gas 1.In other words, there must be 8 electrons (bonded or non-bonded) around atom 1.Non-bonded electron-pair 1.Called “lone pair”

47. Let’s do some examples on the board H 2 –Duet rule F 2 –Octet rule O 2 N 2

48. Lewis structures Example Write the Lewis Structure for the following molecules: 1)H 2 O 2)CCl 4 1)Where does the carbon go & why? 3)PH 3 4)H 2 Se 5)C 2 H 6

49. Lewis structures continued 6)CO 2 7)C 2 H 4 8)C 2 H 2 9)SiO 2

50. Polyatomic ions  If positive charge on ion  Take away electron from central species  If negative charge on ion  Add electron to central species  Example:  H 3 O +

51. Your turn NH 4 + ClO - OH -

52. Electronegativity Electronegativity : The relative ability of an atom in a molecule to attract shared electrons to itself. Example: Fluorine has the highest electronegativity. Similar electronegativity's between elements give non- polar covalent bonds (0.0-0.4) Different electronegativity's between elements give polar covalent bonds (0.5-1.9) If the difference between the electronegativity's of two elements is about 2.0 or greater, the bond is ionic

53. Electronegativity Example For each of the following pairs of bonds, choose the bond that will be more polar.  Al-P vs. Al-N  C-O vs. C-S

54. Dipole moment  Dipole Moment  A molecule that has a center of positive charge and a center of negative charge  Will line up on electric field  In Debye units  1 D = 3.34 x 10 -30 C  m

55. Examples F 2 CO 2 H 2 O NH 3 BF 3 CCl 4

56. Formula Weights Formula weight is the sum of the atomic masses. Example - CO 2 Mass: C + O + O 12.011 + 15.994 + 15.994 43.999

57. Practice Compute the mass of the following compounds round to nearest tenth & state type of bond: NaCl; 23 + 35 = 58; Ionic Bond C 2 H 6 ; 24 + 6 = 30; Covalent Bond Na(CO 3 ) 2 ; 23 + 2(12 + 3x16) = 123; Ionic & Covalent