1. COVALENT BONDING

2. 8.1 Molecules & Molecular Compounds Molecule: a neutral group of atoms joined by covalent bonds Diatomic Molecule: two atoms joined by a covalent bond Examples: H 2, Cl 2, O 2, NO, CO Diatomic elements: Dr. Brinclhof Molecular Compounds: Compounds composed of molecules (covalent bonds)

3. Comparison of Molecular & Ionic Compounds MolecularIonic BondingCovalentIonic Melting pointLowerHigher Electrolyte Usually weak or non Strong Physical state @ room temp (s), (l), (g)(s)

4. Molecular Formulas Show number & type of atoms in a molecule CH 4 H 2 S HNO 3 C 6 H 6 C 3 H 7 OH (NH 4 ) 3 PO 4

5. Structural Formulas Show the arrangement of atoms in a molecule

6. 8.2 Nature of Covalent Bonding Octet rule is a guide Some exceptions will occur Boron accepts less than an octet Phosphorus & Sulfur can accept more than an octet “expanded octet” Electron pairs are shared to form a covalent bond In most cases, octets are completed by sharing pairs of electrons

7. Formation of a Single Covalent Bond Formed when two atoms share one pair of electrons

8. Why do some elements form diatomic molecules?

9. Single Covalent Bonds The hydrogen and oxygen atoms attain noble-gas configurations by sharing electrons.

10. Ammonia, NH 3

11. Drawing Electron Dot (Lewis) Structures Lewis structure is a type of structural formula that depicts all the valence electrons in the molecule or ion See Tutorial 1. Determine the total # ve 2. Connect atoms in such a way that all have a noble gas configuration (octet rule) 3. Carbon is often a central atom 4. Check

12. Draw Lewis Structures for these Molecular Compounds HCl hydrogen chloride Cl 2 chlorine I 2 iodine H 2 O 2 hydrogen peroxide PCl 3 phosphorous trichloride CH 4 methane

13. Single, Double and Triple Covalent Bonds Sometimes atoms share more than one pair of ve’s A bond that involves one shared pair of e - s is a single covalent bond Two shared pairs of electrons is a double covalent bond. Three shared pairs of electrons is a triple covalent bond.

15. Acetylene A gas used in cutting steel Molecular formula is C 2 H 2 Draw the Lewis structure for acetylene 1. Connect the atoms 2. Calculate ve’s 3. Form single covalent bonds between atoms 4. Complete octets until remainder of ve’s are used 5. Form double or triple bonds if needed to complete octets.

16. Polyatomic Ions Same process except… Add or subtract e - s to account for the charge of the ion, for example [NH 4 ] + ammonium ion [SO 4 ] 2- sulfate ion [ClO] - hypochlorite ion [ClO 2 ] - chlorite ion

17. Coordinate Covalent Bonds Bonds in which one of the shared pair comes completely from one of the bonding atoms Carbon Monoxide

18. Bond Energies Energy required to break a chemical bond Energy released when a bond is formed Is a measure of the strength of the bond Large bond energies = strong bonds Type of bond Bond Energy (kJ/mol) C─C347 C=C657 C≡C908

19. Resonance Resonance occurs when two or more valid Lewis structures are possible for a compound or ion Often occurs with placement of a double bond about a central atom Resonance structures are all the valid structures The actual structure is a hybrid of all the possible resonance structures i.e. the bonding present in the particle is a hybrid of those shown in the resonance structures

20. Ozone Is an allotropic form of oxygen Molecular formula is O 3 Is a pollutant (smog) Protects earth by absorbing UV radiation Draw the resonant Lewis structures for ozone

21. Nitrogen Dioxide Formed by lightning strikes Molecular formula NO 2 Also a pollutant in automobile exhaust Draw the Lewis structures for NO 2 Why is this an exception to the octet rule?

22. Exceptions to Octet Rule When there is an odd number of ve, NO 2 Less than an octet: Boron BF 3 More than an octet: Phosphorous PCl 5 Sulfur SF 6 Unfilled d-shells accept additional electrons, creating an “expanded” octet

23. 8.3 Bonding Theories Molecular orbitals When covalent bonds form, atomic orbitals merge to form molecular orbitals

24. Sigma and Pi Bonds Sigma form when atomic orbitals merge along the axis between nuclei (internuclear axis) Pi bonds result when atomic orbitals merge to surround the internuclear axis

25. Sigma Bonds σ bonds are present in single covalent bonds.

26. Sigma bond: p-orbital overlap

27. Pi Bonds π bonds are present in double and triple covalent bonds

28. Sigma and Pi Bonds C 2 H 2

29. VSEPR Theory Valence Shell Electron Pair Repulsion Theory The big idea: Because covalent bonds and non-bonding pairs of electrons are areas of negative charge, they repel one another Covalent bonds and non-bonding electrons are called “electron domains”

30. VSEPR Predicts the shape of small molecules According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible. How to predict the shape of the following molecules: 1. Draw the Lewis structure 2. Count the electron domains around the central atom 3. Determine the domain geometry 4. Determine the molecular geometry (the way the atoms are arranged

31. Methane, CH4 Tetrahedron, bond angles of 109.5°

32. Ammonia, NH 3 Trigonal pyramid, 107° Why is this not trigonal planar? Why is the H-N-H bond angle not 109.5 ° ?

33. Water, H 2 O Draw the Lewis structure Determine the total domains Determine the bonding domains Determine the shape of the molecule Why is water a bent molecule and not a linear one?

34. Molecular Shape of Water Water is a bent molecule Triatomic bent

35. Hybrid Orbitals When covalent bonds form, atomic orbitals mix together to form hybrid orbitals Atomic orbitals involved in bonding often contain a single unpaired electron When the orbitals hybridize, a pair of electrons is shared These hybrid orbitals are equal in number to the atomic orbitals which made them

36. Covalent Bond formation in CH 4 In order for carbon’s 4 ve to be used in bonding, one 2s 2 electron is promoted to 2p. This results in 4 unpaired ve, which can then bond with unpaired e’s of other atoms. In order to accomplish this, the atomic orbitals of C containing these ve hybridize. One s and three p orbitals hybridize to form four equivalent orbitals, called sp 3 orbitals

37. Covalent bonding in CH 4 The s (one) and p (three) orbitals in the valence shell of C hybridize (merge) to form four equivalent sp 3 orbitals. They are called sp 3 orbitals because they are formed from one s orbital and three p orbitals

38. Formation of Hybrid Orbitals http://www.mhhe.com/physsci/chemistry/essentialchemistr y/flash/hybrv18.swf http://www.mhhe.com/physsci/chemistry/essentialchemistr y/flash/hybrv18.swf

39. Hybrid Orbitals Hybridization Involving Single Bonds

40. Hybrid Orbitals Hybridization Involving Double Bonds

41. Hybrid Orbitals Hybridization Involving Triple Bonds

42. How to Determine Hybridization about an Atom The principle: the number of hybrid orbitals must equal the number of atomic orbitals hybridized Count the number of covalent bonds about an atom This must equal the number of hybridized orbitals Beginning with s, continue to add orbitals until the total equals the number of covalent bonds about the atom

43. Hybridization Chart # bondsHybridization 2sp 3sp 2 4sp 3 5?? 6

44. Predicting Hybridization What hybridzation would be found about carbon in the following molecules? HC≡CH sp H 2 C=CH 2 sp 2 H 3 C-CH 3 sp 3

45. 8.4 Polar Bonds and Molecules Electrons in a covalent bond are attracted to the nuclei of both atoms. Why?

46. Unequal Sharing of Bonding Electrons When covalently bonded to another atom, some atoms attract electrons more strongly than others These atoms have greater “electronegativity” When bonded atoms differ in electronegativity, they do not share the bonding electrons equally

47. Bonding Electrons in HCl Bonding e’s spend more time near Cl than H What does this imply about Cl? What does this imply about the distribution of electrical charge in HCl?

48. Polar Covalent Bonds When bonded atoms are sufficiently different in electronegativity, the bond develops negative (-) and positive (+) ends Why? Because the bonding e ’s spend more time around the more electronegative element i.e. the bonding e ’s are not shared equally This unequal distribution of (-) charge is called a dipole The bond is called a polar covalent bond

49. Polar Bonds and Molecules Bond Polarity Bond polarity has to do with unequal distribution of shared electrons caused by differences in electronegativity between bonded atoms This causes one end of the bond to have a “partial positive” (δ + ) charge and the other to have a “partial negative” (δ - )charge These polar covalent bonds and possess a dipole moment The dipole moment is symbolized as -|-------->

50. Bond Character Describes the type of charge distribution in a chemical bond Based upon differences in electronegativity

51. Differences in Electronegativity and Bond Character

52. Polar Molecules Molecules containing polar bonds may have a net dipole The molecule may have a (+) and (-) side Depends upon two factors Presence of polar bonds Geometry (shape) of molecule

53. Polarity of Molecules A molecule as a whole has a dipole depending upon The presence of polar bond(s) The geometry of a molecule Examples: CH 4 CO 2 H 2 O

54. Intermolecular Forces Types of intermolecular forces account for differences between ionic and molecular substances.

55. Polar Molecules

56. Intermolecular Forces of Attraction Not chemical bonds Much weaker than covalent or ionic bonds Van der Waals Forces dipole-dipole interactions London dispersion forces Hydrogen Bonds very important

57. Hydrogen Bonds Hydrogen bonds Attraction between a hydrogen covalently bonded to a very electronegative atom to an unshared electron pair of another electronegative atom May involve different molecules or occur within very large molecules like proteins or nucleic acids http://www.chem.ucla.edu/harding/IGOC/H/hydrogen_bond_acceptor.html

58. Hydrogen Bonding Hydrogen bonding accounts for the unusual properties of water.

59. Hydrogen Bonding in Water http://www.mikeblaber.org/oldwine/BCH4053/Lecture03/Lecture03.htm