1. Unit 7 Covalent Bonding

2. Bonding A metal & a nonmetal transfer electrons –An ionic bond Two metals mix –An alloy (Metallic bond) What do two nonmetals do? –Neither one will give away an electron –So they share their valence electrons –This is a covalent bond

3. Covalent Bonding Nonmetals hold on to their valence electrons They can’t give away electrons to bond Still want to be stable! –Need noble gas configuration (octet rule) Get it by sharing valence electrons with each other. By sharing, both atoms get to count the electrons toward noble gas configuration.

4. Covalent Bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l Both end with full orbitals FF

5. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons… l Both end with full orbitals FF 8 Valence electrons

6. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF 8 Valence electrons

7. Ways to Illustrate Covalent Bonds Molecular formula: shows the number of atoms of each element in a molecule. –Ex. PF 3 Lewis Structures: uses dots to represent bonding between molecular compounds Structural Formulas: shows the arrangement of atoms and bonds –Shared electron dots are replaced with a dash Models: ball and stick (3-D versions)

8. Single Covalent Bond Occurs between nonmetals or a nonmetal & hydrogen Sharing of two valence electrons (1 pair) Different from an ionic bond – electrons are SHARED not transferred

9. An example with dots… It’s like a jigsaw puzzle You will be given the formula You put the pieces together to make everyone stable or happy –Most atoms need an octet –H & He need a duo –Carbon is often the center –Save H for last

10. Water H O Each hydrogen has 1 valence electron and wants 1 more The oxygen has 6 valence electrons and wants 2 more They share to make each other “happy”

11. Water Put the pieces together The first hydrogen is happy The oxygen still wants one more H O

12. Water The second hydrogen attaches Every atom has full energy levels H O H

13. Structural formula… Replace shared dots with a dash O H H

14. Practice – Dots & Structures H 2 CH 4 H 2 S NH 3 CH 2 Cl 2 C 2 H 6 SCl 2 AsF 3 SiH 4 CHF 3

15. Multiple Bonds Sometimes atoms share more than one pair of valence electrons. A double bond is when atoms share two pair (4) of electrons. A triple bond is when atoms share three pair (6) of electrons.

16. Carbon dioxide CO 2 - Carbon is central atom Carbon has 4 valence electrons Wants 4 more Oxygen has 6 valence electrons Wants 2 more O C

17. Carbon dioxide Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short O C

18. Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

19. Carbon dioxide l The only solution is to share more O C O

20. Carbon dioxide l The only solution is to share more O C O

21. Carbon dioxide l The only solution is to share more O CO

22. Carbon dioxide l The only solution is to share more O CO

23. Carbon dioxide l The only solution is to share more O CO

24. Carbon dioxide l The only solution is to share more O CO

25. Carbon dioxide l The only solution is to share more l CO 2 requires two double bonds l Each atom gets to count all the atoms in the bond O CO

26. Carbon dioxide l The only solution is to share more l CO 2 requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

27. Carbon dioxide l The only solution is to share more l CO 2 requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

28. Carbon dioxide l The only solution is to share more l CO 2 requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

29. Carbon Dioxide Replace the shared pairs with dashes O CO

30. Practice O 2 CS 2 CH 2 O N 2 F 2 NO 2 HCN (triple) C 2 H 2 (triple)

31. Exceptions to the Octet Rule/Patterns of Bonding 1. Some elements with odd number of valence electrons –BF 3 –PCl 5 2. Coordinate covalent bonding

32. Coordinate Covalent Bond When one atom donates both electrons in a covalent bond Carbon monoxide (CO) OC

33. Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide (CO) OC

34. Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide (CO) OC OC

35. Summary of Covalent Bonding Covalent bonds occur by SHARING electrons Occurs between NONMETALS End product is called a MOLECULE 1. Molecular compound - formed with different elements 2. Diatomic molecules - 2 of the same atom –There are 7 elements that always form diatomic molecules –H 2, N 2, O 2, F 2, Cl 2, Br 2, and I 2

36. Diatomic Molecules

37. Naming Molecular Compounds Easier than ionic compounds No balancing charges 1 mono- 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa – 7 hepta – 8 octa – 9 nona – 10 deca –

38. Naming Molecular Compounds 1 st element – add the prefix that matches the subscript –Exception – do not add “mono-” if there is 1 atom –No aa, oo, or ao double vowels 2 nd element – add the prefix that matches the subscript –Still ends in -ide

39. Naming CO 2 –1 st element = Carbon; subscript 1 Remember exception “Carbon” –2 nd element = oxygen; subscript 2 Prefix for 2 is di- “Dioxide” –Full name = carbon dioxide

40. Practice Naming S 2 Cl 2 –Disulfur dichloride CS 2 –Carbon disulfide SO 3 –Sulfur trioxide P 4 O 10 –Tetraphosphorus decoxide

41. Name  Formula Just look at the prefixes! Carbon tetrachloride –1 Carbon, 4 Chlorine atoms –CCl 4 Iodine heptaflouride Dinitrogen monoxide Sulfur dioxide

42. Common Names H 2 O – dihydrogen monoxide –Water NH 3 – Nitrogen trihydride –Ammonia CH 4 – carbon tetrahydride Methane HCl – Hydrogen monochloride –Hydrochloric acid

43. Molecular Shapes Lewis diagrams & structural formulas are 2-dimensional Real molecules are 3-D If there are 2 atoms, the molecule has a LINEAR shape (no other options!) –Carbon monoxide (CO) If it has more than 2, how do we figure out the shape?

44. VSEPR Theory Valence Shell Electron Pair Repulsion Theory Used to predict shape of a molecule Negative electrons repel each other and pairs want to be as far apart as possible

45. Linear Linear: 2 atoms around central atom, no unshared pairs on central atom With three atoms the farthest the two outer molecules can get apart is 180º. Will require 2 double bonds or one triple bond C O O 180º

46. Trigonal Planar Trigonal planar: 3 atoms around central atom, no unshared pair on central atom. Angle = 120°

47. 4 molecules around a central atom All single bonds Must think in 3-D! CHH H H Tetrahedral

48. Tetrahedral: 4 atoms around central atom, no unshared pair on central atom A pyramid with a triangular base. C HH H H 109.5º

49. So far… SHAPE# ATOMS FROM THE CENTRAL ATOM # UNSHARED PAIRS ON THE CENTRAL ATOM LINEAR 20 TRIGONAL PLANAR 30 TETRAHEDRAL 40

50. Molecular Shapes But what if there are unshared pairs on the central atom? They still repel each other…

51. Bent OH H O H H <109.5º Bent: 2 atoms around central atom, 1 or 2 unshared pair(s) on central atom.

52. Bent Ball and stick model does not show unshared electron pairs

53. Pyramidal NHH H N HH H <109.5º Pyramidal: 3 atoms around central atom, 1 unshared pair on central atom

54. Pyrimidial Ball and stick model does not show unshared electron pairs

55. Trigonal Bipyramidal Trigonal bipyramidal: 5 atoms around central atom, no unshared pair on central atom Angles = 90° and 120°

56. Adding to the chart… SHAPE# ATOMS FROM THE CENTRAL ATOM # UNSHARED PAIRS ON THE CENTRAL ATOM BENT 21 or 2 PYRAMIDIAL 31 TRIGONAL BIPYRAMIDAL 50

57. So how do I determine the shape of a given molecule? 1.Draw the Lewis diagram 2.Count the atoms and unshared pairs off the central atom 3.Use the VSEPR Theory to determine the shape

58. Which type of bond is it? Also called ionic character Look at which elements are involved –Metal & nonmetal = ionic bond –2 nonmetals = covalent bond Electronegativity – measure of a tendency of an atom to attract a pair of electrons –Influenced by amount of positive charge in the nucleus & electron shielding

59. Differences in Electronegativity. Big difference between values (greater than 1.70) –One atom REALLY wants the electrons and the other…not so much –Ionic bonding  ionic compound Smaller difference between values (less than 1.70) –Both have “equal” attraction for the e - –Covalent bonding  molecule

61. Differences in Electronegativity Medium difference –Still a bit of a tug of war over e- –Unequal sharing of electrons –Results in a POLAR covalent bond –Positive and negative poles –Dipole – partially negative on one side, partially positive on the other

62. Polar Covalent Bond

64. Differences in Electronegativity Very small difference –Share electrons equally –NONPOLAR covalent bond –No positive and negative poles

65. Nonpolar Covalent Bond

67. Polar vs. nonpolar molecules Look at polarity of each bond –All nonpolar bonds = nonpolar molecule (O 2 ) Use the structural formula to draw the molecule & look at the overall shape –Symmetrical polar bonds cancel each other out so molecule = nonpolar (CO 2, CCl 4 ) –Nonsymmetrical polar bonds = polar molecule (H 2 O)

68. Dipole-dipole attraction Attraction between + part of one dipole and - part of another dipole Hydrogen bond - between an electronegative atom and a hydrogen atom bonded to another electronegative atom –Often involves F, N, or O –Strongest of the intermolecular forces

69. Hydrogen Bonding

70. Surface Tension/H Bonding

71. Van der Waals – London dispersion force Weak intermolecular force caused by negative electrons on one side of a cloud being attracted to a nearby positive nucleus Constantly changing

72. Properties of Molecular Compounds Poor conductors of heat & electricity Often found as liquids or gases Weaker attraction between atoms Low melting & boiling points

73. IONIC vs COVALENT

74. Carbon (Organic) Chemistry Carbon plays a dominant role in the chemistry of living things Bonding stability –4 valence electrons –Very unlikely to form ionic bonds –Can form covalent bonds with LOTS of different elements (especially H & O) Small molecules link together resulting in the formation of a large variety of structures often with repeating subunits

75. Examples of carbon-based compounds Simple hydrocarbons Small carbon molecules with functional groups Complex polymers (including biological molecules)

76. Simple Hydrocarbons Petrochemicals – Propane, Butane, Octane Structure Notation

77. Hydrocarbons “Saturation” with hydrogens affects shape & reactivity –Saturated vs. unsaturated fats – shape –Reactivity based on energy needed to maintain double and triple bonds

78. Alkane vs Alkene vs Alkyne

79. Easy to identify from condensed structural formulas Butane – CH 3 -CH 2 -CH 2 -CH 3 Butene – CH 3 -CH 2 -CH-CH 2 Butyne – CH-C-CH 2 -CH 3

80. Alkane vs Alkene vs Alkyne Also easy to identify from full structural fomulas

81. Functional groups Specific groups of atoms that are responsible for chemical characteristics of a compound ALWAYS a close relationship between properties & structure (aspirin, vitamins, insulin)

82. Complex Polymers Natural polymers –Proteins, nucleic acids Synthetic polymers –Polythene, Polystyrene –Kevlar –Nylon

83. Isomers Molecules with the same chemical formula but a different structure

84. Allotropy Property of some elements that allows them to exist in different forms based on their connectivity Graphene –Charge your phoneCharge your phone

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