1. 1-1 William H. Brown Beloit College William H. Brown Christopher S. Foote Brent L. Iverson Eric Anslyn http://academic.cengage.com/chemistry/brown Chapter 1 Covalent Bonding and Shapes of Molecules

2. 1-2 Organic Chemistry  The study of the compounds of carbon.  Over 10 million compounds have been identified. About 1000 new ones are identified each day!  C is a small atom. It forms single, double and triple bonds. It is intermediate in electronegativity (2.5). It forms strong bonds with C, H, O, N, and some metals.

3. 1-3 Schematic View of an Atom A small dense nucleus, diameter 10 -14 - 10 -15 m, which contains positively charged protons and most of the mass of the atom. An extranuclear space, diameter 10 -10 m, which contains negatively charged electrons.

4. 1-4 The Concept of Energy  In the ground state of carbon, electrons are placed in accordance with the quantum chemistry principles (aufbau, Hund’s rule, Pauli exclusion principle, etc.) that dictate the lowest energy form of carbon.  If we place the electrons in a different manner (as for example with one electron in the 2s and three electrons in the 2p) we would have a higher energy level referred to as an excited state. When the electrons are rearranged back to the ground state, energy is released.

5. 1-5 The Concept of Energy  Electrons in the lowest energy orbital, 1s, are held tightest to the nucleus and are the hardest to remove from the atom.  First ionization energy: The energy needed to remove the most loosely held electron from an atom or molecule.

6. 1-6 Shapes of Atomic s and p Orbitals All s orbitals have the shape of a sphere with the center of the sphere at the nucleus. Figure 1.8 (a) Calculated and (b) cartoon representations showing an arbitrary boundary surface containing about 95% of the electron density.

7. 1-7 Shapes of Atomic s and p Orbitals Figure 1.9 (a) Three-dimensional representations of the 2p x, 2p y, and 2p z atomic orbitals computed using the Schrödinger equation. Nodal planes are shaded.

8. 1-8 Shapes of Atomic s and p Orbitals  Figure 1.9(b) Cartoon representations of the 2p x, 2p y, and 2p z atomic orbitals.

9. 1-9 Molecular Orbital Theory  MO theory begins with the hypothesis that electrons in atoms exist in atomic orbitals and electrons in molecules exist in molecular orbitals.

10. 1-10 Molecular Orbital Theory  Figure 1.10 MOs derived from combination by (a) addition and (b) subtraction of two 1s atomic orbitals.

11. 1-11 VB: Hybridization of Atomic Orbitals The number of hybrid orbitals formed is equal to the number of atomic orbitals combined. Elements of the 2nd period form three types of hybrid orbitals, designated sp 3, sp 2, and sp. The mathematical combination of one 2s atomic orbital and three 2p atomic orbitals forms four equivalent sp 3 hybrid orbitals.

12. 1-12 VB: Hybridization of Atomic Orbitals  Figure 1.12 sp 3 Hybrid orbitals. (a) Computed and (b) cartoon three-dimensional representations. (c) Four balloons of similar size and shape tied together, will assume a tetrahedral geometry.

13. 1-13 VB: Hybridization of Atomic Orbitals  Figure 1.13 Orbital overlap pictures of methane, ammonia, and water.

14. 1-14 VB: Hybridization of Atomic Orbitals  Figure 1.14 sp 2 Hybrid orbitals and a single 2p orbital on an sp 2 hybridized atom.

15. 1-15 VB: Hybridization of Atomic Orbitals  The mathematical combination of one 2s atomic orbital and one 2p atomic orbital gives two equivalent sp hybrid orbitals.

16. 1-16  Figure 1.16 sp Hybrid orbitals and two 2p orbitals on an sp hybridized atom. VB: Hybridization of Atomic Orbitals

17. 1-17 Combining VB & MO Theories  A double bond uses sp 2 hybridization.  Consider ethylene, C 2 H 4. Carbon and other second-period elements use a combination of sp 2 hybrid orbitals and the unhybridized 2p orbital to form double bonds.

18. 1-18 Combining VB & MO Theories  A carbon-carbon triple bond consists of one  bond formed by overlap of sp hybrid orbitals and two  bonds formed by the overlap of parallel 2p atomic orbitals.

19. 1-19 Covalent Bonding of Carbon

20. 1-20 Polar and Nonpolar Molecules  To determine if a molecule is polar, we need to determine if the molecule has polar bonds and the arrangement of these bonds in space.  Molecular dipole moment (  ):  Molecular dipole moment (  ): The vector sum of the individual bond dipole moments in a molecule. reported in Debyes (D)

21. 1-21 Electrostatic Potential (elpot) Maps  Relative electron density distribution in molecules is important because it allows us to identify sites of chemical reactivity. Many reactions involve an area of relatively high electron density on one molecule reacting with an area of relatively low electron density on another molecule. It is convenient to keep track of overall molecular electron density distributions using computer graphics.

22. 1-22 Electrostatic Potential (elpot) Maps  In electrostatic potential maps (elpots) Areas of relatively high calculated electron density are shown in red. Areas of relatively low calculated electron density are shown in blue. Intermediate electron densities are represented by intermediate colors.

23. 1-23 Polar and Nonpolar Molecules  These molecules have polar bonds, but each molecule has a zero dipole moment.

24. 1-24 Polar and Nonpolar Molecules  These molecules have polar bonds and are polar molecules.

25. 1-25 Polar and Nonpolar Molecules Formaldehyde has polar bonds and is a polar molecule.

26. 1-26 Resonance  For many molecules and ions, no single Lewis structure provides a truly accurate representation.

27. 1-27 Resonance  Examples: equivalent contributing structures.

28. 1-28 Resonance  Curved arrow:  Curved arrow: A symbol used to show the redistribution of valence electrons.  In using curved arrows, there are only two allowed types of electron redistribution: from a bond to an adjacent atom. from a lone pair on an atom to an adjacent bond.  Electron pushing is a survival skill in organic chemistry. learn it well!

29. 1-29 Resonance  All contributing structures must 1. have the same number of valence electrons. 2. obey the rules of covalent bonding: no more than 2 electrons in the valence shell of H. no more than 8 electrons in the valence shell of a 2nd period element. 3. differ only in distribution of valence electrons; the position of all nuclei must be the same. 4. have the same number of paired and unpaired electrons.

30. 1-30 Resonance  The carbonate ion Is a hybrid of three equivalent contributing structures. The negative charge is distributed equally among the three oxygens as shown in the elpot.

31. 1-31 Resonance  Preference 4:  Preference 4: negative charge on the more electronegative atom. Structures that carry a negative charge on the more electronegative atom contribute more than those with the negative charge on a less electronegative atom.

32. 1-32 Lewis Dot Structures  Gilbert N. Lewis  Valence shell: The outermost occupied electron shell of an atom.  Valence electrons: Electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions.  Lewis dot structure: The symbol of an element represents the nucleus and all inner shell electrons. Dots represent electrons in the valence shell of the atom.

33. 1-33 Lewis Dot Structures  Table 1.4 Lewis Dot Structures for Elements 1-18

34. 1-34 Lewis Model of Bonding  Atoms interact in such a way that each participating atom acquires an electron configuration that is the same as that of the noble gas nearest it in atomic number. anionAn atom that gains electrons becomes an anion. cationAn atom that loses electrons becomes a cation. ionic solids. This ionic interaction is often referred to as an ionic bondThe attraction of anions and cations leads to the formation of ionic solids. This ionic interaction is often referred to as an ionic bond. covalent bond polar covalent bondsAn atom may share electrons with one or more atoms to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bond. Bonds may be partially ionic or partially covalent; these bonds are called polar covalent bonds

35. 1-35 Electronegativity  Electronegativity: A measure of an atom’s attraction for the electrons it shares with another atom in a chemical bond.  Pauling scale Generally increases left to right in a row. Generally increases bottom to top in a column.

36. 1-36 Covalent Bonds  The simplest covalent bond is that in H 2 The single electrons from each atom combine to form an electron pair. The shared pair functions in two ways simultaneously; it is shared by the two atoms and fills the valence shell of each atom.  The number of shared pairs One shared pair forms a single bond Two shared pairs form a double bond Three shared pairs form a triple bond

37. 1-37 Polar and Nonpolar Covalent Bonds  Although all covalent bonds involve sharing of electrons, they differ widely in the degree of sharing.  We divide covalent bonds into nonpolar covalent bonds and polar covalent bonds.

38. 1-38 Polar and Nonpolar Covalent Bonds An example of a polar covalent bond is that of H-Cl. The difference in electronegativity between Cl and H is 3.0 - 2.1 = 0.9.  +  -We show polarity by using the symbols  + and  -, or by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end.

39. 1-39 Polar Covalent Bonds  Bond dipole moment (  ): A measure of the polarity of a covalent bond. The product of the charge on either atom of a polar bond times the distance between the two nuclei. Table 1.7 shows average bond dipole moments of selected covalent bonds.

40. 1-40 Lewis Structures  To write a Lewis structure Determine the number of valence electrons. Determine the arrangement of atoms. Connect the atoms by single bonds. Arrange the remaining electrons so that each atom has a complete valence shell. Show a bonding pair of electrons as a single line. Show a nonbonding pair of electrons (a lone pair) as a pair of dots. In a single bond atoms share one pair of electrons, in a double bond they share two pairs of electrons and in a triple bond they share three pairs of electrons.

41. 1-41 Lewis Structures - Table 1.8  In neutral molecules hydrogen has one bond. carbon has 4 bonds and no lone pairs. nitrogen has 3 bonds and 1 lone pair. oxygen has 2 bonds and 2 lone pairs. halogens have 1 bond and 3 lone pairs.

42. 1-42 Formal Charge  Formal charge:  Formal charge: The charge on an atom in a molecule or a polyatomic ion.  To derive formal charge 1. Write a correct Lewis structure for the molecule or ion. 2. Assign each atom all its unshared (nonbonding) electrons and one-half its shared (bonding) electrons. 3. Compare this number with the number of valence electrons in the neutral, unbonded atom. 4. The sum of all formal charges is equal to the total charge on the molecule or ion.

43. 1-43 Apparent Exceptions to the Octet Rule  Molecules that contain atoms of Group 3A elements, particularly boron and aluminum.

44. 1-44 Apparent Exceptions to the Octet Rule  Atoms of third-period elements, such as S and P, are often drawn with more bonds than allowed by the octet rule. The P in trimethylphosphine obeys the octet rule by having three bonds and one unshared pair. A common depiction of phosphoric acid, however, has five bonds to P, which is explained by invoking the use of 3d orbitals to accommodate the additional bonds.

45. 1-45 Apparent Exceptions to the Octet Rule  However, the use of 3d orbitals for bonding is in debate.  An alternative representation that gives P in phosphoric acid an octet has four bonds and a positive formal charge on P. The oxygen involved in the double bond of the alternative depiction has one bond and a negative formal charge.

46. 1-46 Apparent Exceptions to the Octet Rule Sulfur, another third-period element, is commonly depicted with varying numbers of bonds. In each of the alternative structures sulfur obeys the octet rule, and has one or more positive formal charges.

47. 1-47 Functional Groups  Functional group:  Functional group: An atom or group of atoms within a molecule that shows a characteristic set of physical and chemical properties.  Functional groups are important for three reason; they are: 1. the units by which we divide organic compounds into classes. 2. the sites of characteristic chemical reactions. 3. the basis for naming organic compounds.

48. 1-48 Alcohols hydroxyl  Contain an -OH (hydroxyl) group bonded to a tetrahedral carbon atom. condensed structural formula.  Ethanol may also be written as a condensed structural formula.

49. 1-49 Alcohols Alcohols are classified as primary (1°), secondary (2°), or tertiary (3°) depending on the number of carbon atoms bonded to the carbon bearing the -OH group.

50. 1-50 Alcohols There are two alcohols with molecular formula C 3 H 8 O.

51. 1-51 Amines amino group  Contain an amino group; an sp 3 -hybridized nitrogen bonded to one, two, or three carbon atoms. An amine may by 1°, 2°, or 3°.

52. 1-52 Aldehydes and Ketones carbonyl (C=O) group.  Contain a carbonyl (C=O) group.

53. 1-53 Carboxylic Acids carboxyl (-COOH) group.  Contain a carboxyl (-COOH) group.

54. 1-54 Carboxylic Esters  Ester:  Ester: A derivative of a carboxylic acid in which the carboxyl hydrogen is replaced by a carbon group.

55. 1-55 Carboxylic Amide  Carboxylic amide amide  Carboxylic amide, commonly referred to as an amide: A derivative of a carboxylic acid in which the -OH of the -COOH group is replaced by an amine. The six atoms of the amide functional group lie in a plane with bond angles of approximately 120°.

56. 1-56 Covalent Bonds & Shapes of Molecules End Chapter 1