1. Unit 6: Chemical Bonding Refer to Ch. 8 & 9 for supplemental reading.

2. 1. Chemical Bond: an attractive force that holds 2 atoms together 3 types: ionic, covalent, metallic 2. Valence Electrons: the electrons in the outer energy level of an atom. They are like the front lines of an army. They are the electrons involved in bonding. Review

3. Review How do you find valence electrons? How do you find valence electrons? Hint there are two ways! Hint there are two ways! Examples: Examples: Mg ___ Mg ___ O ___ O ___ Ar ___ Ar ___ Si ___ Si ___ Examples: Examples: Mg Mg O 2 6 8 4 1s 2 2s 2 2p 6 3s 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2p 1s 2 2s 2 2p 4

4. Electron Dot Structures Depicts element symbol w/ valence e - shown as dots. Depicts element symbol w/ valence e - shown as dots. Na MgAl ClAr Si ON

5. Ionic Bonds Occurs when ions of opposite charge (+,-) attract each other. Occurs when ions of opposite charge (+,-) attract each other. Metal ion – Nonmetal ion Metal ion – Nonmetal ion Simplest attraction Simplest attraction NaCl MgF 2 NaCl MgF 2 Polyatomic ions Polyatomic ions AlPO 4 (NH 4 ) 2 SO 4 AlPO 4 (NH 4 ) 2 SO 4

6. Formation of Ionic Bond Cation- positive ion (+) Cation- positive ion (+) Forms when a metal atom loses e - to become stable. Forms when a metal atom loses e - to become stable. Anion- negative ion (-) Anion- negative ion (-) Forms when a nonmetal atom gains e - to become stable Forms when a nonmetal atom gains e - to become stable An ionic bond is formed when e - are transferred between atoms and the resulting ions stick together. An ionic bond is formed when e - are transferred between atoms and the resulting ions stick together.

7. Examples Formation of NaCl Formation of NaCl Na Cl → [Na] + [ Cl ] - = NaCl Na Cl → [Na] + [ Cl ] - = NaCl Formation of MgF 2 Formation of MgF 2 Mg F F  [Mg] 2+ [ F ] - [ F ] - = MgF 2 Mg F F  [Mg] 2+ [ F ] - [ F ] - = MgF 2 How would a compound form between two aluminum and three oxygen? How would a compound form between two aluminum and three oxygen?

8. Electron Configuration of Ions Cation example: (metal) Cation example: (metal) Ca atom:1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Ca atom:1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Ca 2+ ion:1s 2 2s 2 2p 6 3s 2 3p 6 Lost 2 electrons to obtain noble gas configuration (octet) Ca 2+ ion:1s 2 2s 2 2p 6 3s 2 3p 6 Lost 2 electrons to obtain noble gas configuration (octet)

9. Electron Configuration of Ions Anion example: (nonmetal) Anion example: (nonmetal) N atom:1s 2 2s 2 2p 3 N atom:1s 2 2s 2 2p 3 N 3- ion:1s 2 2s 2 2p 6 N 3- ion:1s 2 2s 2 2p 6 Gained 3 electrons to obtain noble gas configuration (octet)

10. Properties of Ionic Compounds IONIC Bond Formation Type of Structure Boiling Point Electrical Conductivity Other Properties high yes * (solution or liquid) high Melting Point Crystal lattice Physical State Solid (hard and rigid) brittle e - transferred from metal to nonmetal

11. Electrolyte A substance that conducts electricity A substance that conducts electricity Because of ionic bonds ionic (charged) nature, ionic compounds conduct electricity in the molten or aqueous forms. Because of ionic bonds ionic (charged) nature, ionic compounds conduct electricity in the molten or aqueous forms.

12. Nonpolar Nonpolar Polar Polar Ionic Ionic View Bonding Atomic Bonding : Chemistry ZoneAnimations.Atomic Bonding : Chemistry Zone C. Bond Polarity

13. Covalent Bonds Occurs when 2 nonmetals share pairs of electrons to become stable. Molecular compounds are formed. Occurs when 2 nonmetals share pairs of electrons to become stable. Molecular compounds are formed. Examples: Examples: H 2 OCO 2 C 6 H 12 O 6 PCl 5 H 2 OCO 2 C 6 H 12 O 6 PCl 5

14. Covalent Bonds Covalent bonds can be Covalent bonds can be single (1 shared pair) single (1 shared pair) double (2 shared pairs) double (2 shared pairs) or triple (3 shared pairs) or triple (3 shared pairs) Bond strength: triple > double > single Bond strength: triple > double > single Bond length: single > double > triple Bond length: single > double > triple

15. 2 4 8888 8 6 8 NEED

16. 1 2 4567 8 3 Figure out through e config Available

17. Creating Lewis Structures Follow this system: Follow this system: Example: H 2 O Example: H 2 O 1) Draw a “skeleton” of the molecule. It generally works to place the “different” atom in the center. 1) Draw a “skeleton” of the molecule. It generally works to place the “different” atom in the center. H O H

18. Creating Lewis Structures Find the needed electrons (N) for each atom and add them up. N will be 8 for most elements, with these exceptions: Find the needed electrons (N) for each atom and add them up. N will be 8 for most elements, with these exceptions: H gets 2 valence e - H gets 2 valence e - Be gets 4 valence e - Be gets 4 valence e - B gets 6 valence e - B gets 6 valence e - N = 12 H = 2 O = 8 H = 2

19. 3) Find the available (valence) electrons (A) for each atom and then add them up*. 3) Find the available (valence) electrons (A) for each atom and then add them up*. A = A = *special note: when completing a Lewis structure for a polyatomic ion, you will need to correct A by adding the absolute value of the charge if negative, and subtracting the charge if positive. For example, for the ion PO 4 3-, you would add 3 to A. For the ion NH 4+, you would subtract 1 from A. (You do the opposite of the charge.) *special note: when completing a Lewis structure for a polyatomic ion, you will need to correct A by adding the absolute value of the charge if negative, and subtracting the charge if positive. For example, for the ion PO 4 3-, you would add 3 to A. For the ion NH 4+, you would subtract 1 from A. (You do the opposite of the charge.) H = 1 O = 6 H = 1 Total A = 8 N = 12 A = 8

20. 4) Find the shared (S) electrons for the entire molecule by this formula: S = N – A 4) Find the shared (S) electrons for the entire molecule by this formula: S = N – A S = S = S= 12 – 8 = 4 N = 12 A = 8 S = 4

21. 5) The shared electrons are the bonding electrons. Place all of the shared electrons between the atoms. 5) The shared electrons are the bonding electrons. Place all of the shared electrons between the atoms. H O H 6) You must place all of the available (A) electrons in the picture. The shared electrons are part of the available. See how many of the available electrons still need to be placed, and put them in the picture as lone pairs (unshared pairs) so that every atom gets an octet (remember H only needs 2). 6) You must place all of the available (A) electrons in the picture. The shared electrons are part of the available. See how many of the available electrons still need to be placed, and put them in the picture as lone pairs (unshared pairs) so that every atom gets an octet (remember H only needs 2). H O H N = 12 A = 8 S = 4 N = 12 A = 8 S = 4 4

22. H O H

23. F B F F

24. F S F F

25. CF 4 CF 4 N=A=S= F F C F F 8+(4x8) = 40 4+(4x7) = 32 8 24

26. BeCl 2 BeCl 2 N=A=S= 4+(2x8) = 20 2+(2x7) = 16 4 12 Cl Be Cl

27. CO 2 CO 2 N=A=S= 8+(2x8) = 24 4+(2x6) = 16 8 8 O C O

28. Polyatomic Ions To find total # of valence e - (A): To find total # of valence e - (A): Add 1e - for each negative charge. Add 1e - for each negative charge. Subtract 1e - for each positive charge. Subtract 1e - for each positive charge. Place brackets around the ion and label the charge. Place brackets around the ion and label the charge.

29. Polyatomic Ions ClO 4 - ClO 4 - O O Cl O O N=A=S= 8+(4x8) = 40 7+(4x6) = 31 24 +1 =32 8

30. Bond Polarity Most bonds are a blend of ionic and covalent characteristics. Most bonds are a blend of ionic and covalent characteristics. Difference in electronegativity determines bond type. Difference in electronegativity determines bond type. If  EN is: Bond type is: < 0.4Nonpolar covalent  EN < 1.7 Polar covalent > 1.7Ionic

31. Nonpolar Covalent Bond Nonpolar Covalent Bond e - are shared equally e - are shared equally symmetrical e - density symmetrical e - density usually identical atoms usually identical atoms Bond Polarity

32. ++ -- Polar Covalent Bond Polar Covalent Bond e - are shared unequally e - are shared unequally asymmetrical e - density asymmetrical e - density results in partial charges (dipole) results in partial charges (dipole)

33. If ΔEN is:Bond type is: < 0.4Nonpolar covalent 0.4 < Δ EN < 1.7Polar covalent > 1.7Ionic Examples: Cl 2 Cl 2 HCl HCl NaCl NaCl 3.16-3.16=0.0Nonpolar2.2-3.16=0.96Polar.93-3.16=2.23Ionic

34. VSEPR Theory Valence Shell Electron Pair Repulsion Theory Valence Shell Electron Pair Repulsion Theory Electron pairs orient themselves in order to minimize repulsive forces. Electron pairs orient themselves in order to minimize repulsive forces.

35. A. VSEPR Theory Types of e - Pairs Types of e - Pairs Bonding pairs - form bonds Bonding pairs - form bonds Lone pairs - nonbonding e - Lone pairs - nonbonding e - Lone pairs repel more strongly than bonding pairs!!!

36. A. VSEPR Theory Lone pairs reduce the bond angle between atoms. Lone pairs reduce the bond angle between atoms. Bond Angle

37. Draw the Lewis Diagram. Draw the Lewis Diagram. Tally up e - pairs on central atom. Tally up e - pairs on central atom. double/triple bonds = ONE pair double/triple bonds = ONE pair Shape is determined by the # of bonding pairs and lone pairs. Shape is determined by the # of bonding pairs and lone pairs. Know the 8 common shapes & their bond angles! B. Determining Molecular Shape

38. Common Molecular Shapes 2 total 2 bond 0 lone LINEAR180° BeH 2

39. Common Molecular Shapes 3 total 2 bond 1 lone BENT<120° SO 2

40. 4 total 2 bond 2 lone BENT109.5° H2OH2OH2OH2O Common Molecular Shapes

41. 3 total 3 bond 0 lone TRIGONAL PLANAR 120° BF 3 Common Molecular Shapes

42. 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° NH 3 Common Molecular Shapes

43. 4 total 4 bond 0 lone TETRAHEDRAL109.5° CH 4 Common Molecular Shapes

44. PF 3 PF 3 4 total 3 bond 1 lone PYRAMIDAL107° F P F F Examples Examples

45. CO 2 CO 2 O C O 2 total 2 bond 0 lone LINEAR180°Examples

46. Examples H 2 SCCl 4 BF 3 SiO 2