Covalent Bonding

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: H2, Cl2, O2, NO, CO • Diatomic elements: Dr. Brinclhof • Molecular Compounds: Compounds composed of molecules (covalent bonds)

3. Comparison of Molecular & Ionic Compounds

4. Molecular Formulas • Show number & type of atoms in a molecule • CH4, H2S • HNO3 • C6H6 • C3H7OH • (NH4)3PO4

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

6. 8.2 Nature of Covalent Bonding • Octet rule is a guide • Electrons are sharedto form a covalent bond

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, NH3

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 • Determine the total # ve • Connect atoms in such a way that all have a noble gas configuration (octet rule) • Carbon is often a central atom • Check

12. Draw Lewis Structures for these Molecular Compounds • HCl hydrogen chloride • Cl2 chlorine • I2 iodine • H2O2 hydrogen peroxide • PCl3 phosphorous trichloride • CH4 methane

13. Single, Double and Triple Covalent Bonds • Sometimes atoms share more than one pair of ve’s • A bond that involves on 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 C2H2 • Draw the Lewis structure for acetylene • Connect the atoms • Calculate ve’s • Form single covalent bonds between atoms • Complete octets until remainder of ve’s are used • 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 • [NH4]+ • [SO4]2-

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

19. Resonance Structures • A resonance structure is a structure that occurs when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion. • Actual bonding is a hybrid of all the possible resonance structures

20. Ozone • Is an allotropic form of oxygen • Molecular formula is O3 • 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 NO2 • Also a pollutant in automobile exhaust • Draw the resonance structures for NO2 • Why is this an exception to the octet rule?

22. Exceptions to Octet Rule • When there is an odd number of ve, NO2 • Less than an octet: • BoronBF3 • More than an octet: • PhosphorousPCl5 • SulfurSF6 • 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 bonds result atomic orbitals merge along the axis between nuclei (internuclear axis) • Pi bonds result when atomic orbitals merge to around the internuclear axis

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

26. Pi Bonds πbonds are present in double and triple covalent bonds

27. Sigma and Pi BondsC2H2

28. 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”

29. 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: • Draw the Lewis structure • Count the electron domains • Determine the geometry of the molecule (the way the atoms are arranged

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

31. Ammonia, NH3 Trigonal pyramid, 107° Why is this not trigonal planar? Why is the H-N-H bond angle not 109.5 °?

32. Water, H2O • Draw the Lewis structure • Determine the total domains • Determine the bonding domains • Determine the shape of the molecule • Why is water a bend molecule and not a linear one?

33. 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

34. Covalent Bond formation in CH4 In order for carbon’s 4 ve to be used in bonding, one 2s2electron 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 sp3 orbitals

35. Covalent bonding in CH4 • The s (one) and p (three) orbitals in the valence shell of C hybridize (merge) to form four equivalent sp3 orbitals. • They are called sp3 orbitals because they are formed from one s orbital and three p orbitals

36. Formation of Hybrid Orbitals • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/hybrv18.swf

37. Hybrid Orbitals • Hybridization Involving Single Bonds

38. Hybrid Orbitals • Hybridization Involving Double Bonds

39. Hybrid Orbitals • Hybridization Involving Triple Bonds

40. 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

41. Hybridization Chart

42. Predicting Hybridization • What hybridzation would be found about carbon in the following molecules? • HC≡CH • sp • H2C=CH2 • sp2 • H3C-CH3 • sp3

43. 8.4 Polar Bonds and Molecules • Electrons in a covalent bond are attracted to the nuclei of both atoms. Why?

44. 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

45. 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?

46. 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 • This unequal distribution of (-) charge is called a dipole • The bond is called a polar covalent bond

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

48. Differences in Electronegativity and Bond Character

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

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