1. Chapter 1 Lecture Introduction and Review Organic Chemistry, 8 th Edition L. G. Wade, Jr.

2. Lecture 1: Atoms, Chemical Bonding, and Drawing Lewis Structures Atoms & Isotopes Electronic Structure of Atom Electron Configurations Ionic & Covalent Bonding Electronegativity Dipole Moment Lewis Structures 2

3. Chapter 13 Organic Chemistry Organic chemistry is the chemistry of carbon compounds.

5. Chapter 95 Atoms & Isotopes Two Structures ◦ Nucleus: holds protons (p + ) and neutrons (n 0 ) ◦ Very small, highly dense (+) central core # p + = Z = Atomic # Chapter 15

6. Chapter 96 Atoms & Isotopes Two Structures ◦ Electron Cloud: houses electrons (e - ) ◦ Very large volume; negatively charged exterior of atom # p + = # e - for any neutral atom Chapter 16

7. Chapter 97 Atoms & Isotopes Atoms with the same #p + but different #n 0 result in isotopes of the same element. Isotopes identified by Mass Number (A) A = # p + + # n 0 Chapter 97

8. Atoms & Isotopes Concept ReviewEASY ◦ 1. Write an alternate definition for an isotope that relates mass number and atomic number. ◦ 2. Determine the number of protons, neutrons, and electrons in a 17 N isotope. 8

9. Electronic Structure of Atom Protons/Neutrons are ~1800 times heavier than electrons Electrons have both particle and wave properties; called Wave- Particle Duality Electrons: affect bonding / reactivity / chemical properties Located outside nucleus in orbitals ◦ Allowed energy states for an electron Heisenberg Uncertainty Principle – can never exactly determine position and momentum of an electron 9

10. Electronic Structure of Atom: Orbitals An orbital is described by a wave function. ◦ Helps describe distribution of electron density in space. ◦ Squaring the wave function gives the electron density 10

11. Electronic Structure of Atom: Orbitals Electron density is the probability of finding the electron in a particular part of an orbital. Every electron is characterized by a set of quantum numbers. Orbitals are grouped into “shells” at different distances from the nucleus. ◦ Shells are identified by principal quantum number (n) ◦ Depicts energy level for electron with valuesn = {1, 2, 3, …∞} ◦ As n increases, shells are farther from the nucleus. 11

12. Electronic Structure of Atom: Orbitals Within each shell are “sub-shells” that hold electrons. ◦ Sub-shells are defined by azimuthal quantum number (l) ◦ Depicts shape of orbital:l = {0,1, 2, 3, …(n-1)} ◦ Sub-shell names:s-orbitals spherical (l = 0) p-orbitalsdumb-bell (l = 1) d-orbitalsfour-leaf clover (l = 2) The third quantum number depicts the orientation in three-dimensional space for each orbital. ◦ Magnetic quantum number m l :m l = {-l, …, 0, …+l} ◦ Values for m l reveal how many orbitals are in each sub-shell 12

13. Electronic Structure of Atom: Orbitals The fourth quantum number is called the “spin quantum number” and is denoted by the symbol m s. ◦ Shows how electrons enter each individual orbital ◦ Only two values possible:m s = { 1/2, -1/2} Pauli Exclusion Principle states that each orbital holds two electrons with opposing spins. 13

14. First Electron Shell The 1s orbital holds two electrons. 14

15. Chapter 915 Second Electron Shell 2s orbital (spherical) Three 2p orbitals (dumb-bell) Chapter 915

16. Electronic Structure of Atom Concept Review:MEDIUM ◦ Derive the set of quantum numbers (n, l, m l and m s ) that would characterize the last electron added to a neutral oxygen atom? Explain what each quantum number describes about the electronic structure of this electron. 16

17. Electronic Configurations Aufbau principle states that electrons fill the lowest energy orbitals first. Hund’s rule states that degenerate (equal energy) orbitals will fill with one electron before pairing occurs. ◦ Electrons repel each other, so energy is needed to couple pair Chapter 117

18. Electronic Configurations 18

19. Electronic Configurations Valence electrons are electrons on the outermost shell of the atom. Chapter 119

20. Electronic Configurations Concept Review:EASY 1. Draw the orbital energy diagram for a phosphorus atom. How many core electrons and valence electrons are present? 2. Write out the correct electron configuration for a calcium atom. 20

21. Ionic Bonding To obtain a noble gas configuration (a full valence shell), atoms may transfer electrons from one atom to another. The atoms, now bearing opposite charges, attract each other, forming an ionic bond. Chapter 121

22. Covalent Bonding Electrons are shared between the atoms to complete the octet. When the electrons are shared evenly, the bond is said to be nonpolar covalent, or pure covalent. When electrons are not shared evenly between the atoms, the resulting bond will be polar covalent. Chapter 122

23. Electronegativity and Bond Polarity Greater ΔEN means greater polarity. Chapter 123

24. Polar Covalent Bonds Electrons tend to shift away from lower electronegativity atoms to higher electronegativity atoms. The greater the difference in electronegativity, the more polar the bond. 24

25. Dipole Moment Dipole moment is defined to be the amount of charge separation (  ) multiplied by the bond length (  ). Charge separation is shown by an electrostatic potential map (EPM), where red indicates a partially negative region and blue indicates a partially positive region. Chapter 125

26. Lewis Structures For simple Lewis structures… 1.Calculate total number of valence electrons in the structure.  Draw individual atoms with dots showing valence electrons 2.Put the atom’s together so they share PAIRS of electrons to make complete octets.  Atoms that share an electron pair between them are called bonding electrons. 3.Extra electrons are placed on atoms starting from outside and working toward the center. Electrons found isolated on an atom are called nonbonding or lone pair electrons. 4.Check total number valence electrons to verify right sketch. 26

27. Lewis Structures A First Example ◦ Draw the Lewis Structure for MethaneCH 4 27

28. Lewis Structures A dash represents a bonding pair of electrons. Chapter 128

29. Lewis Structures Some more simple examples ◦ Draw the Lewis structure for H 2 O ◦ Draw the Lewis structure for NH 3 Now Try This: ◦ Draw the Lewis structure for ethane C 2 H 6 29

31. Lewis Structures Concept Review:DIFFICULT ◦ Draw Lewis structures for the following. Circle all lone pair electrons in your diagram(s).  CH 5 N  C 2 H 5 OH  C 3 H 7 Cl 31

32. Lewis Structures How do we deal with an ion (charged atom)? 1.For an anion, we add one electron for each negative charge to the total number of valence electrons. 2.For a cation, we subtract one electron for each positive charge. Key Points to Remember ◦ In drawing Lewis structures we try to give each atom the electron configuration of a noble gas. (Octet Rule) ◦ Certain elements do not follow Octet Rule; one must learn to recognize these exceptions Chapter 132

33. Exceptions to Octet Rule 1.Hydrogen and Helium never follow the Octet Rule; instead they follow the Duet Rule (two electrons in first shell) 2.Elements in the 2 nd row of periodic table usually obey the octet rule; an atom that resides in second row can never exceed an octet (eight electrons in second shell). 3.Elements in 3 rd row of periodic table may not obey the octet rule. Chapter 133

34. Common Bonding Patterns Neutral atoms in organic molecules should follow the pattern shown above. Valence in this case refers to the number of covalent bonds that atom can make; carbon is tetravalent (makes 4 bonds); oxygen is divalent (makes 2 bonds). 34

35. Multiple Bonding Sharing two pairs of electrons is called a double bond. Chapter 135

36. Multiple Bonding Sharing three pairs of electrons is called a triple bond. Chapter 136

37. Lewis structures are the way we write organic chemistry. Learning now to draw them quickly and correctly will help you throughout this course. Chapter 137

38. Skill Building: Practice Problems Problems 1-1 thru 1-5 38