Unit 6: Theories of Covalent Bonding and Intro. To Organic Chemistry

1. Unit 6: Theories of Covalent Bonding and Intro. To Organic Chemistry • Lewis Structures • VSEPR • Theories of Covalent Bonding • Valence Bond Theory • Molecular Orbital Theory • Organic Chemistry • Functional Groups • Nomenclature • Simple Reactions

2. Chemical Bonds • Octet Rule: Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons. • When ionic compounds are formed, electrons are gained or lost. • When molecular compounds are formed, electrons are shared.

3. Chemical Bonds • Chemical bond:strong attractive force that exists between atoms (or ions) in a compound • ionic bonds • covalent bonds • metallic bonds

4. Chemical Bonds • Ionic Bond: the electrostatic force of attraction between oppositely charged ions in an ionic compound • metal cation (+) • non-metal anion (-) • The Na+ and Cl- ions in a salt (NaCl) crystal are held together by electrostatic attraction.

5. Chemical Bonds • Covalent Bonds: the attractive force between atoms in a molecule that results from sharing of one or more pairs of electrons • non-metals • H2O : • Cl2 : H-O and Cl-Cl bonds result from sharing of electrons

6. Lewis Symbols • Valence electronsare involved in chemical bonding: • electrons residing in the incomplete outer shell of an atom • For main group elements, the number of valence electrons for an element = group number of the element • N (group 5A) has 5 valence electrons • Br (group 7A) has 7 valence electrons

7. Lewis Symbols • Lewis symbols (electron-dot symbols) are used to depict valence electrons in an atom or ion • chemical symbol for the element • dot for each valence electron • dots are placed on all 4 sides of the chemical symbol • all four sides of the symbol are equivalent • up to 2 dots (electrons) per side

8. Lewis Symbols Example: Draw the Lewis symbol for oxygen. Example: Draw the Lewis symbol for carbon.

9. Covalent Bonding • Lewis structures (also called electron-dot structures) can be used to represent the covalent bonds that are present in a molecule. • The formation of H2: H + H  H H or H H

10. Covalent Bonding • Components of Lewis (electron-dot) structures: • Elemental symbol for each atom • Bond between atoms depicted using a solid line • Unshared electron pairs are shown around the appropriate atom

11. Cl Cl O C O N N Covalent Bonding • Single bond: • one pair of shared electrons • Double bond • Two pairs of shared electrons • Triple bond • Three pairs of shared electrons

12. Drawing Lewis Structures To draw a Lewis structure: • Add up the valence electrons from all atoms • For a cation (+), subtract 1 electron for each positive charge • NH4+ : 5 + 4 (1) -1 = 8 e- • For an anion (-), add 1 electron for each negative charge • CN-: 4 + 5 + 1 = 10 e-

13. Drawing Lewis Structures • Write the chemical symbols for each atom showing which is attached to which using a single bond (-). • Sometimes (but not always) the order in which the formula is written • HCN: H-CN • Central atom (often written first) surrounded by other atoms • Commonly the least electronegative element (except H) will be the central atom • CCl4 : C with 4 Cl attached to it

14. Drawing Lewis Structures • Add electron pairs to the atoms bonded to the central atom first until each has an octet of electrons. • Remember, H only gets 2 electrons

15. Drawing Lewis Structures • Place any leftover electrons on the central atom. • Sometimes results in more than an octet on the central atom

16. Drawing Lewis Structures • If there are not enough electrons to give the central atom an octet, try multiple bonds. • Use one (or more) unshared pairs of electrons from an outer atom to form double (or triple) bonds H C N H C N

17. Drawing Lewis Structures Example: Draw the Lewis structure for COCl2.

18. Drawing Lewis Structures Example: Draw the Lewis structure for the carbonate ion.

19. Drawing Lewis Structures Example:Draw all possible resonance structure for CO2.

20. Drawing Lewis Structures Example: Draw all possible resonance structures for SCN-. Which resonance structure is the major contributor to the resonance hybrid?

21. VSEPR • Lewis structures show the number and type of bonds between atoms in a molecule. • All atoms are drawn in the same plane (the paper). • Do not show the shape of the molecule

22. VSEPR • The valence-shell electron-pair repulsion model (VSEPR) can be used to predict the shape of an ABn molecule when A is a main group element. ABn where A = central atom, main group element B = outer atoms n = # of “B” atoms Examples: CO2, H2O, BF3, NH3, CCl4

23. VSEPR • VSEPR counts the number of electron domainsaround the central atom where electrons are likely to be found and uses this number to predict the shape. • Electron domains: regions around the central atom where electrons are likely to be found. • Two types of electron domains are considered: • bonding pairs of electrons • nonbonding (lone) pairs of electrons

24. VSEPR • Bonding pairs of electrons: electrons that are shared between two atoms Cl Cl C Cl Cl Bonding pairs Bondingpairs

25. VSEPR • Nonbonding (lone) pairs of electrons: • electrons that are found principally on one atom • unshared electrons H N H H Nonbonding pair

26. VSEPR • Ammonia (NH3) has 4 electron domains: H N H H 1 nonbonding pair 3 bonding pairs

27. VSEPR • Electron domains tend to repel each other • regions of high electron density • like charges repel each other • According to VSEPR, the best arrangement of a given number of electron domains is the one that minimizes repulsions between them. • Electron domain geometry: • The arrangement of electron domains around the central atom

28. Electron Domain Geometries You must know these! You must be able to draw these!

29. Electron Doman Geometries Trigonal planar Tetrahedral Trigonal bipyramidal octahedral

30. VSEPR • In order to determine the electron domain geometry: • draw the Lewis structure • count the total # of electron domains • multiple bonds = 1 electron domain • determine the electron-domain geometry

31. VSEPR Example: Predict the electron domain geometry of IF5.

32. VSEPR • The electron domain geometry does NOT tell you the actual shape of the molecule. • Molecular geometry: the arrangement of the atoms in space • Molecular geometry is a consequence of electron-domain geometry.

33. VSEPR Example: Identify and draw the molecular geometry of IF5.

34. VSEPR Example: Identify and draw the electron domain and molecular geometries for I3-.

35. Valence Bond Theory • Covalent bonds form when atoms share electrons • Electron density is concentrated between the nuclei • Two common theories are used to explain the properties of molecules in terms of the bonding that exists between atoms. • Valence bond theory • Molecular orbital theory

36. Valence Bond Theory • According to valence bond theory, an electron pair bond is formed between two atoms when a valence atomic orbital on one atomoverlaps with a valence atomic orbital on another atom. • Overlap:share a region of space • 2 electrons of opposite spin share common space between the nuclei

37. Valence Bond Theory • Overlap can occur between two s orbitals, two p orbitals, or one s and one p orbital: • Overlap between two s orbitals overlap

38. Valence Bond Theory • Overlap between one s and one p orbital overlap

39. Valence Bond Theory • Overlap between two p orbitals overlap

40. Valence Bond Theory • The previous bonds are called s bonds. • electron density is concentrated symmetrically along an imaginary line connecting the two nuclei (internuclear axis) Internuclear axis

41. Valence Bond Theory • p bonds are formed when sideways overlap occurs between two p orbitals that are oriented perpendicular to the internuclear axis. • Overlap regions lie both above and below the internuclear axis

42. Valence Bond Theory • p bonds are involved in the formation of double and triple bonds • Single bond: • one s bond • Double bond: • one s bond and one p bond • Triple bond: • one s bond and two p bonds

43. Valence Bond Theory • Sometimes, simple overlap between s and/or p orbitals can’t explain the actual shape or properties of compounds. • Consider a BeF2 molecule: F Be F

44. F F 1s22s22p5 1s22s22p5 Be 1s22s2 Valence Bond Theory One unpaired e- One unpaired e- How can Be form covalent bonds with F if it doesn’t have unpaired electrons??? No unpaired electrons

45. Be Be 1s22s12p1 1s22s2 Valence Bond Theory Promote one of the 2s e- to a 2p orbital no unpaired electrons 2 unpaired electrons

46. F F 1s22s22p5 1s22s22p5 Be 1s22s12p1 Valence Bond Theory One unpaired e- One unpaired e- • Predicted overlaps: • 2s (Be) – 2p (F) • 2p (Be) – 2 p (F) Implies two different kinds of Be – F bonds! 2 unpaired electrons

47. Valence Bond Theory • Both of the Be – F bonds in BeF2 are identical! • The solution: • Mix the Be 2s orbital with one of the Be 2p orbitals to form two hybrid orbitals • atomic orbitals formed by mixing 2 or more atomic orbitals on an atom

48. Valence Bond Theory • Two sp hybrid orbitals are formed when one s and one p orbital are hybridized.

49. F F 1s22s22p5 1s22s22p5 Be 1s sp 2p Valence Bond Theory • When Be forms covalent bonds with two F, each sp hybrid orbital on the Be atom overlaps with a p orbital located on a F atom.

50. Valence Bond Theory • The BeF2 molecule is linear. • sp hybridization implies that the electron domain geometry around the central atom is linear.