2. REVIEW We can tell how many electrons and atom will gain or lose by looking at its valence. Metals like to lose electrons. (Cations) –Ex. Na + Nonmetals like to gain electrons. (Anions) –Ex: O 2- All elements try to have a full valence of 8 electrons(OCTET RULE).

3. REVIEW Cation- is a positively charged ion. How do cations form? –When atoms LOSE electrons they become positive. Anion- is a negatively charged ion. How do anions form? –When atoms GAIN electrons they become negative.

4. Chemical Bonding Notes A chemical bond is the force of attraction that holds two atoms together. Attractive Force NaCl

5. http://www.wisc- online.com/objects/ViewObject.aspx?ID=G CH2204http://www.wisc- online.com/objects/ViewObject.aspx?ID=G CH2204

6. Why do elements form chemical bonds? Atoms form chemical bonds in order to fill their outermost energy level with electrons. A full valence shell causes an atom to be more stable.A full valence shell causes an atom to be more stable. A full valence shell consists of 8 valence electrons.

7. Ionic Bonding Ionic bonds: Metal atoms transfer electrons to nonmetal atoms. Producing oppositely charged ions (cation & anion) which attract each other. Na + Cl  Na + Cl -

8. Remember: Non-metal atoms take electrons from metal atoms to form an octet.

9. How to write a formula. Write cation first, followed by anion Example: Anion : P 3- Cation : Al 3+ Formula : AlP

10. How to write a formula. Compound must be neutral, so all charges must cancel

11. Write an ionic formula for Na + bonding with F − Balance the charges. Na + F − (1+) + (1-) = 0 1 Na + and 1 F − = NaF

12. Write an ionic formula for Mg 2+ bonding with Cl − Balance the charges. Mg 2+ Cl − Cl − (2+) + 2(1-) = 0 1 Mg 2+ and 2 Cl − = Mg Cl 2

13. Write an ionic formula for K + bonding with S 2− Balance the charges. K + S 2− K + 2(1+) + (2-) = 0 2 K + and 1 S 2− = K 2 S

14. Write the formula for… an ionic compound composed of: Al 3+ and S 2- Al 2 S 3

15. Write an ionic formula for Fe 3+ bonding with OH − Balance the charges. Fe 3+ OH − OH − (3+) + 3(1-) = 0 1 Fe 3+ and 3 OH − = Fe(OH) 3

16. Let’s play the Ionic Bonding Dating Game!

17. Example: Aluminum Chloride Step 1: Step 2: Step 3: 13 Step 4: AlCl 3 Criss-Cross Rule Al Cl 3+ 1- write out name with space write symbols & charge of elements criss-cross charges as subsrcipts combine as formula unit (“1” is never shown) Aluminum Chloride

18. Example: Aluminum Oxide Step 1: Aluminum Oxide Step 2: Al 3+ O 2- Step 3: Al O 23 Step 4: Al 2 O 3 Criss-Cross Rule

19. Example: Magnesium Oxide Step 1: Magnesium Oxide Step 2: Mg 2+ O 2- Step 3: Mg O 22 Step 4: Mg 2 O 2 Step 5: MgO Criss-Cross Rule

20. InBr 3 BaS Criss-Cross Rule criss-cross rule: charge on cation / anion “becomes” subscript of anion / cation ** Warning: Reduce to lowest terms. Al 2 O 3 Al 3+ and O 2– Al 2 O 3 Ba 2+ and S 2– Ba 2 S 2 In 3+ and Br 1– In 1 Br 3 aluminum oxidebarium sulfideindium bromide

21. Lesson Three--Transition Metal Compounds Transition metals have electrons in d orbitals and can donate different numbers of electrons, thus giving them several different positive charges. These can be determined from the Roman numeral which is written next to the metal's name. Example: Cu 1+ is Copper I Pb 2+ is Lead II Fe 3+ is Iron III Sn 4+ s Tin IV

22. Metals with more than 1 charge Examples: Cu + Copper (I) Cu +2 Copper (II) Fe +2 Iron (II) Fe +3 Iron (III)

23. Practice K and Cl K and S Ca and S Cu (II) and S

24. Polyatomic Ions!!!!!!!! A polyatomic ion is a charged species (ion) composed of two or more atoms covalently bonded.ion covalently bonded PO 4 -3 NH 4 +1

25. Lewis Dot Structures for Polyatomic ions H+ NH 4 + H N H H

26. Al + PO 4 -3 K + SO 4 -2 Al +3 + PO 4 -3 K +1 + SO 4 -2 Al(PO 4 )K 2 (SO 4 )

27. Ca + PO 4 -3 Ca +2 + PO 4 -3 Ca 3 (PO 4 ) 2

28. Multiple Oxidation States “tin fluoride” tin (II) fluoridetin (IV) fluoride Tin is either 2+ or 4+ oxidation state. Sn 2+ F 1- SnF 2 Sn 4+ F 1- SnF 4 tin (II) sulfide Sn 2+ S 2- Sn 2 S 2 SnS tin (II) sulfate Sn 2+ SO 4 2- SnSO 4 tin (II) sulfite Sn 2+ SO 3 2- SnSO 3 tin (IV) sulfate Sn 4+ SO 4 2- Sn 2 (SO 4 ) 4 Sn(SO 4 ) 2

30. Covalent bonding Notes Covalent bond: The sharing of a pair of electrons between 2 nonmetal atoms in order to fill its valence shell. –Each atom gains 1 electron from each covalent bond it forms with another atom.

31. When electron sharing usually occurs so that atoms attain a stable electron configuration and have 8 valence electrons.

32. Single Covalent Bonds Diatomic Molecules Each chlorine needs to gain one electron by sharing electrons each atom achieves stability. Cl + Cl  ClCl The pair of shared electrons is often represented as a dash. Cl-Cl

33. Single Covalent Bonds Diatomic Molecules The chlorine atoms only share one pair of valence electrons. The electrons pairs not shared are called unshared electron pairs or lone pairs. Cl + Cl  ClCl

34. Single Covalent Bonds in compounds H 2 0 is a molecule containing three atoms with two single covalent bonds. Count up the electrons you have!!! 2 H + O H O H The hydrogen and oxygen attain stable configurations by sharing electrons.

35. Your Turn Example OF 2

36. Double Covalent Bonds Two pair of electrons are being shared. S + S  S S

37. Triple Covalent Bonds Three pair of electrons are being shared. P + P  P P

38. Charged Compounds Some compounds do not satisfy their stable configuration and therefore have a charge on the compound. Example- NH 4 +1

40. Exceptions to the Octet Rule The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. However, these molecules do exist in nature. Examples: Nitrogen dioxide (NO 2 ) Boron trifluoride (BF 3 ) Phosphorus pentachloride (PCl 5 ) = 10 v.e - Expanded octet Sulfur hexafluroride (SF 6 )= 12 v.e - Expanded octet

42. Nonpolar Covalent Bond When atoms bond equally it is considered a nonpolar covalent bond. Cl 2 O 2 N 2 H 2

43. Polar Covalent Bond When electrons are shared unequally it is a polar covalent bond. An atom that strongly attracts electrons is more electronegative and therefore gains a slightly negative charge. The less electronegative atom has a slightly positive charge. This results in a polar bond!

44. An arrow is used to show which element is donating the unshared pair of electrons. The crossed end of the arrow indicates a pos. end and the arrow points in the direction of the neg. end Example: H-Br

45. polar molecules are also called dipoles. A dipole is a molecule with two partially charged ends or poles.

46. Examples H-Br H 2 S SCl 2 CO 2

47. C. Johannesson Nonpolar Covalent Bond –e - are shared equally –symmetrical e - density –usually identical atoms C. Bond Polarity

48. C. Johannesson ++ -- C. Bond Polarity Polar Covalent Bond –e - are shared unequally –asymmetrical e - density –results in partial charges (dipole)

49. C. Johannesson ++ -- ++ B. Lewis Structures Nonpolar Covalent - no charges Polar Covalent - partial charges

50. C. Johannesson A. Dipole Moment Direction of the polar bond in a molecule. Arrow points toward the more e - neg atom. H Cl ++ --

51. C. Johannesson B. Determining Molecular Polarity Nonpolar Molecules –Dipole moments are symmetrical and cancel out. BF 3 F F F B

52. C. Johannesson B. Determining Molecular Polarity Polar Molecules –Dipole moments are asymmetrical and don’t cancel. net dipole moment H2OH2O H H O

53. C. Johannesson CHCl 3 H Cl B. Determining Molecular Polarity Therefore, polar molecules have... –asymmetrical shape (lone pairs) or –asymmetrical atoms net dipole moment