LET S FIRST REVIEW IONIC BONDING

Transcription

1 COVALENT BONDING

2 LET S FIRST REVIEW IONIC BONDING

3 In an IONIC bond, electrons are lost or gained, resulting in the formation of IONS in ionic compounds. K F

4 K F

5 K F

6 K F

7 K F

8 K F

9 K + F _ The compound potassium fluoride consists of potassium (K + ) ions and fluoride (F - ) ions

10 Ionic Bonds are/have: 1. Metals and Nonmetals 2. Solids 3. High melting points 4. Large differences in electronegativity

11 COVALENT BONDS Ms. Thompson

12 Covalent Bonding tug of war Share electrons Molecules of 2 or more atoms 2 nonmetal molecules Groups 4A, 5A, 6A, and 7A

13 Covalent Molecules Properties Two nonmetals Low melting points Low conductivity Nonelectrolytes

14 Prefixes for Covalent Molecules Prefix Number Mono- 1 Di- 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6 Hepta- 7 Octa- 8 Nona- 9 Deca- 10

15 Naming Molecular Compounds First nonmetal is written as: PREFIX + element name If there is one atom, do NOT use mono Second nonmetal is written as: PREFIX + element name root + ide If there is one atom, USE mono Example: IF 5 Iodine Pentafluoride H 2 O Dihydrogen Monoxide SBr 6 Sulfur Hexabromide

16 Examples CO CO 2 SO 2 CF 4 Carbon monoxide Carbon dioxide Sulfur dioxide Carbon tetrafluoride

17 So what are covalent bonds?

18 In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule).

19 In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair.

20 In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair. The shared electron pair is called a bonding pair

21 Types of Covalent Bonds Single: two atoms share one pair of electrons Ex. H H Double: two shared pairs of electrons Ex. O = O Triple: three shared pairs of electrons Ex. N = N

22 Chlorine forms a covalent bond with itself Cl 2

23 Cl Cl How will two chlorine atoms react?

24 Cl Cl Each chlorine atom wants to gain one electron to achieve an octet

25 Cl Cl Neither atom will give up an electron chlorine is highly electronegative. What s the solution what can they do to achieve an octet?

26 Cl Cl

27 Cl Cl

28 Cl Cl

29 Cl Cl

30 Cl Cl octet

31 Cl Cl octet circle the electrons for each atom that completes their octets

32 Cl Cl The octet is achieved by each atom sharing the electron pair in the middle circle the electrons for each atom that completes their octets

33 Cl Cl The octet is achieved by each atom sharing the electron pair in the middle circle the electrons for each atom that completes their octets

34 Cl Cl This is the bonding pair circle the electrons for each atom that completes their octets

35 Cl Cl It is a single bonding pair circle the electrons for each atom that completes their octets

36 Cl Cl It is called a SINGLE BOND circle the electrons for each atom that completes their octets

37 O2 Oxygen is also one of the diatomic molecules

38 O O How will two oxygen atoms bond?

39 O O Each atom has two unpaired electrons

40 O O

41 O O

42 O O

43 O O

44 O O

45 O O

46 O O Oxygen atoms are highly electronegative. So both atoms want to gain two electrons.

47 O O Oxygen atoms are highly electronegative. So both atoms want to gain two electrons.

48 O O

49 O O

50 O O

51 O O

52 O O Both electron pairs are shared.

53 O O 6 valence electrons plus 2 shared electrons = full octet

54 O O 6 valence electrons plus 2 shared electrons = full octet

55 O O two bonding pairs, making a double bond

56 O O O=O For convenience, the double bond can be shown as two dashes.

57 O=O This is the oxygen molecule, O2

58 Drawing the Lewis Structure 1. Sum the total number of valence electrons 2. The atom usually written first in the chemical formula is the central atom in the Lewis structure 3. Complete the octet bonded to the central atom 4. If there are not enough electrons to give the central atom an octet try multiple bonds

59 VSEPR Valence Shell Electron Pair Repulsion Theory States the shape of molecules is based upon the concept that electrons, being of like charge, will repel themselves to the greatest possible distances. Electron pairs repel Shows three-dimensional shape Electron pairs are as far apart as possible Produces molecular geometry of shapes

60 Shapes of Molecules Linear - No non-bonding electrons Linear Molecules have a bond angle = 180 AB 2 must be either linear or bent: Examples of Linear molecules

61 Linear Molecules have a bond angle = 180 Bent molecules have a bond angle 180 H 2 O AB 2 E or AB 2 E 2 - classification Water is a bent molecule with bond angles of **Notice the bond angle decreases as the number of non-bonding pairs increases**

62 Trigonal Planar B A B AB 3 most common shapes place the B atoms at the corners of an equilateral triangle: The A atom lies in the same plane as the B atoms (Flat) Bond angle = 120 No non-bonding electrons

63 Trigonal Pyramidal The A atom lies above the plane of the B atom. Pyramid with an equilateral triangle as the base.

64 AB 4 is Tetrahedral H H C H H The carbon has 4 valence electrons and thus needs 4 more electrons from four hydrogen atoms to complete its octet. The hydrogen atoms are as far apart as possible at 109 bond angle. This is tetrahedral geometry. The molecule is three dimensional.

65 Six pairs of electrons around the central atom are based on the Octahedron structure. AB 6 The central atom can be visualized as being at the centre of an octahedron, with the six electrons pointing to the six vertices all bond angles are 90 SF 6 BrF 5 XeF 4 Octahedral 90 Square Pyramidal Should be less than 90º Square Planar

66 Further Examples: Tutorial Questions : Draw Lewis structures and the molecular geometry of the following molecules: H 3 O +, NH 4+, CS 2, SCl 2

67 Shape Bondingpairs Nonbonding pairs Bond angle Examples Linear BeCl 2, CO 2, HCN, C 2 H 2 Trigonal planar BF 3, SO 3, NO 3-, CO 2-3, C 2 H 4 Tetrahedral NH 4+, SO 2-4, PO 3-4, Ni(CO) 4, CH 4 Trigonal pyramidal PH 3, SO 3 2-, NH 3 Bent H 2 S, SO 2, H 2 O

68 Predicting the Shape of the Molecule 5. Sum the Number of Electron Domains around the Central Atom in the Lewis Structure; Single = Double = Triple Bonds = Non-Bonding Lone Pair of Electrons = One Electron Domain 6. From the Total Number of Electron Domains, Predict the Geometry and Bond Angle(s); 2 (Linear = 180º); 3 (Trigonal Planar = 120º); 4 (Tetrahedral = 109.5º); 5 (Trigonal Bipyramidal = 120º and 90º); 6 (Octahedral = 90º) 7. Lone Pair Electron Domains exert a greater repulsive force than Bonding Domains. Electron Domains of Multiple Bonds exert a greater repulsive force than Single Bonds. Thus they tend to compress the bond angle.

69 Review Ionic vs. covalent Naming Lewis dot structures (ionic and covalent) Shapes/bond angles

70 Nonpolar vs. Polar Bonds Nonpolar Bond formed between two similar atoms Electrons are shared equally Usually a diatomic molecule or between C & H Polar Bond formed between two atoms of two different elements Electrons are shared unequally Ex. HCl, H 2 O, CO Ex. H 2, N 2, O 2, F 2, Cl 2, I 2, Br 2 (Diatomic Molecules) CH 4, C 2 H 6, C 3 H 8

71 Intermolecular Forces Dispersion Forces (aka van der Waals forces): Weakest Between two nonpolar molecules Usually a diatomic molecule or between C & H Ex. F 2, Cl 2, C 2 H 6, CH 4

72 Intermolecular Forces Dipole-Dipole Forces: Between two polar molecules Slightly negative region is weakly attracted to a slightly positive region Ex. HCl

73 Intermolecular Forces Hydrogen Bonding: A hydrogen bonds to a very electronegative atom (O, N, or F) Ex. H 2 O

74 Ranking based on Polarity Polarity increases with: Increased intermolecular forces Polarity also increases with: Increased boiling point Increased viscosity Increased electronegativity Decreased temperature **Therefore, when ranking molecules by increasing boiling point, you are really ranking the molecules by increasing polarity or intermolecular forces.

75 Types of Intermolecular Forces Practice Wkst Answers 1. Hydrogen 9. Hydrogen 2. Dispersion 10. Dispersion 3. Hydrogen 11. Dipole-Dipole 4. Dispersion 12. Dipole-Dipole 5. Dipole-Dipole 13. Hydrogen 6. Dispersion 14. Dispersion 7. Dispersion 15. Dipole-Dipole 8. Dipole-Dipole

76 Polarity Practice Worksheet Polarity Practice 1. CHBr 3 2. H 2 O 3. HI 4. C 2 HBr 5. CH 3 OH Ranking by Inc Polarity 1. NF 3 < PF 3 < SF 2 < LiOH 2. N 2 H 2 < C 2 H 5 OH, CH 3 OH < Ni(OH) 3 3. B 2 F 4 < H 2 C 2 O 4 < CF 2 O < CuCl 2 4. PH 3 < NF 3 < PF 3 < NH 3 5. H 2 < H 2 S < H 2 O < HF

77 Intermolecular Forces Wkst 1. Dispersion 2. Hydrogen 3. Dispersion 4. Dipole-Dipole 5. Dispersion 6. Dipole-Dipole 7. Dipole- Dipole 8. Hydrogen 11. C 2 H 6 < C 2 H 5 < C 2 H 5 OH 12. H 2 <H 2 S < H 2 O 13. BI 3 < BBr 3 < BCl CH 4 < CH 3 OCH 3 < CH 4 O < CaCO 3

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79 Electronegativity & Nonpolar vs. Polar Bonds Electronegativity Difference Range Type of Bond Example Nonpolar Covalent H H (0.0) Moderately Polar Covalent H Cl (0.9) Very Polar Covalent H F (1.9) 2.0 Ionic Na + Cl - (2.1)