1. Chapter 1 Chemical Bonding and Chemical Structure

2. Organic chemistry 2 The branch of chemistry that deals with carbon based compounds – Organic compounds may contain any number of other elements, including hydrogen, nitrogen, oxygen, halogens, phosphorus, silicon, and sulfur MethaneSucroseMorphine

3. History Vitalism: Only biological systems (e.g., plants, animals) could produce organic compounds Wohler’s synthesis of urea (1828), began to undermine vitalism

4. Why Study Organic Chemistry? Organic chemistry lies at the heart of the modern chemical industry Central to medicine and pharmacy Interface of physical and biological sciences Everyday applications: Plastics, textiles, communications, transportation, food, clothing, cosmetics, etc. 4

5. Review of Chemical Bonding Valence Electrons: Outermost electrons s and p electrons for main group elements Responsible for chemical properties of atoms Participate in chemical reactions Core Electrons Valence Electron

6. Octet Rule Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells – Natural electron configuration of the Noble Gases – Done by gaining, losing, or sharing electrons – Increases stability – H and He seek a “Duet”

7. Ionic Bonding Ions: atoms that have a charge due to gain or loss of electrons – Anion: (-) charged atom – Cation: (+) charged atom Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

8. Formula Unit

11. Ionic bonds are omni-directional Can dissociate into free ions 11

12. Covalent Compounds Covalent Compounds: compounds composed of atoms bonded to each other through the sharing of electrons Electrons NOT transferred No + or – charges on atoms Non-metal + Non-metal Also called “molecules” Examples : – H 2 O – CO 2 – Cl 2 – CH 4

13. or H-H or Duet

14. Covalent Bonds 14

15. Electronegativity The measure of the ability of an atom to attract electrons to itself – Increases across period (left to right) and – Decreases down group (top to bottom) – fluorine is the most electronegative element – francium is the least electronegative element

16. Electronegativity Scale

17. Types of Bonding 1)Non-Polar Covalent Bond: Difference in electronegativity values of atoms is 0.0 – 0.4 Electrons in molecule are equally shared Examples: Cl 2, H 2, CH 4 EN Cl = 3.0 3.0 - 3.0 = 0 Pure Covalent

18. 2)Polar Covalent Bond: Difference in electronegativity values of atoms is 0.4 – 1.7/2.0 Electrons in the molecule are not equally shared The atom with the higher EN value pulls the electron cloud towards itself Partial charges Examples: HCl, ClF, NO EN Cl = 3.0 EN H = 2.1 3.0 – 2.1 = 0.9 Polar Covalent

19. Electrostatic Potential Maps A graphical depiction of electron distribution 19

20. 3)Ionic Bond: Difference in EN above 1.7-2.0 Complete transfer of electron(s) Whole charges EN Cl = 3.0 EN Na = 1.0 3.0 – 0.9 = 2.1 Ionic

22. Dipole Moment (  ) Depends on charge separation and distance  = qr (a vector quantity) q = magnitude of charge r = vector from site of + charge to site of – charge Units = Debyes (D)

23. Molecular Polarity

26. Lewis Dot Structures 1)Count the number of valence electrons present in the molecule 2)Determine the arrangement of atoms. Generally, the atom that occurs least often is central. Join the terminal atoms to the central atom(s) using shared pairs of electrons (bonds) 3)Place any remaining electrons around the terminal atoms to satisfy the octet rule Exception: Hydrogen 4)Place any remaining electrons on the central atom(s) to satisfy the octet rule

27. 5)Check to make sure: You’ve used the correct number of valence electrons Everyone has an octet (or duet) Everyone is doing what they like to do 6)If the number of electrons around the central atom is less than 3, change the single bonds to multiple bonds

28. What Things Like To Do 1)Halogens Like to be terminal Like to have one bonding pair (two shared electrons) and 3 lone pairs (non-bonding electrons) 2)Carbon Likes to have 4 bonding pairs and no lone pairs Likes to bond to other carbons Likes to be central 3)Silicon Likes to do what carbon does Notice, it sits under C on the periodic table

29. 4)Oxygen Like to have 2 bonding pairs and 2 lone pairs 5)Sulfur Likes to do what O does 6)Nitrogen Likes to have 3 bonding pairs and 1 lone pair 7)Phosphorous Likes to do what N does

30. 8)Hydrogen Likes to be terminal with only 1 bond Do not put lone pairs on H 9)Boron Likes to have 3 bonds and no lone pairs Likes a sextet instead of an octet (what everybody else besides Hydrogen likes) 10)*Note: A double bond = 2 bonding pairs A triple bond = 3 bonding pairs

31. Problems Draw the Lewis Dot Structures for the following molecules 1)CO 2 2)P 2 H 4 3)O 3 4)NO 3 -

33. 33 Drawing Resonance Structures 1.Draw first Lewis structure that maximizes octets 2.Assign formal charges 3.Move electron pairs from atoms with (-) formal charge toward atoms with (+) formal charge

34. Formal Charge Assigned charge for each atom in a molecule/ion – Electronic bookkeeping – may or may not correspond to a real charge – Sum of formal charges on each atom must equal the total charge on the molecule/ion FC = Valence e - ’s – Lone Pair e - ’s – ½ bonding e - ’s

35. Molecular Structures of Covalent Compounds Atomic connectivity: How atoms in a molecule are connected Molecular geometry: How far apart atoms are and how they are arranged in space – Bond lengths – Bond angles – Dihedral angles OR

36. Bond Length Distance between nuclei 36

37. Increases with atoms in higher rows Decreases toward higher atomic number along a row Decreases with increasing bond order 37

38. Bond Angles Angle between each pair of bonds Contribute to molecular shape Determined by Valence-shell electron-pair repulsion (VSEPR) Use molecular models! Line-and-wedge structures 38

39. Drawing LDS With Correct Geometry

40. Valence Shell Electron Pair Repulsion Theory VSEPR theory: – Electrons repel each other – Electrons arrange in a molecule themselves so as to be as far apart as possible Minimize repulsion Determines molecular geometry

43. Defining Molecular Shape Electron pair geometry: the geometrical arrangement of electron groups around a central atom – Look at all bonding and non-bonding e - ’s Molecular Geometry: the geometrical arrangement of atoms around a central atom – Ignore lone pair electrons

46. Problems Predict the approximate geometry in each of the following molecules – BF 3 – HCN – CO 3 2- Estimate the bond angles and relative bond lengths in the following molecule

48. Dihedral Angle Also known as the torsional angle Rotation can occur along single bonds 48

49. Valence Bond Theory

56. Types of Bonds A sigma (  ) bond results when the bonding orbitals point along the axis connecting the two bonding nuclei – either standard atomic orbitals or hybrids s-to-s, p-to-p, hybrid-to-hybrid, s-to-hybrid, etc. A pi (  ) bond results when the bonding orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei – between unhybridized parallel p orbitals the interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore  bonds are stronger than  bonds

60. Problems Write a hybridization and bonding scheme for acetaldehyde

61. Molecular Orbital Theory

62. Bond Order: ½ (# of electrons in bonding MO’s - # of electrons in antibonding MO’s)

65. Problems 1)Draw an MO diagram to predict the bond order of N 2 2)Draw an MO diagram to predict the bond order of CN - 3)Use MO theory to determine the bond order of Ne 2