Chapter 1

1. Chapter 1 The Basic – Bonding and Molecular Structure

2. Organic Chemistry and Life

4. 1.1 The Development of Organic Chemistry as Science • Organic compounds: compounds that could be obtained from living organisms • The scientific study of the structure, properties, composition, reactions, and preparation (by synthesis or by other means) of chemical compounds that contain carbon • Inorganic compounds: those came from non-living sources • Occur as a salts

5. Atomic Orbitals

6. Atomic Orbitals S - orbital p- orbital

7. Electrons Configuration • Show

8. Orbital Diagram and Electron Configuration • Electrons configuration: H, He, C, Mg • Aufbau Principle: fill lowest orbital first to full capacity, then next

9. 1.2 The structural Theory of Organic Chemistry • Atoms in organic compounds can form a fixed number of bonds using their valence electrons

10. 1.2 The structural Theory of Organic Chemistry • A carbon atom can use one or more of its valence electrons to form bonds to other carbon atoms

11. 1.3 Isomers: The Importance of Structural Formulas • Constitutional isomers – non identical compounds with same molecular formula • Do not necessary share similar properties

12. 1.4 Ionic Bonds Occurs in ionic compound Results from transferring electron Created a strong attraction among the closely pack compound

13. Covalent Bonding • Formation of a covalent Bond • Two atoms come close together, and electrostatic interactions begin to develop • Two nuclei repel each other; electrons repel each other • Each nucleus attracts to electrons; electrons attract both nuclei • Attractive forces > repulsive forces; then covalent bond is formed

14. Electronegativity • Electronegativity (EN): the ability of an atom in a molecule to attract the shared electron in a bond • Metallic elements – low electronegativities • Halogens and other elements in upper right-hand corner of periodic table – high electronegativity

15. Polarity • Polar covalent bonds – the bonding electrons are attracted somewhat more strongly by one atom in a bond • Electrons are not completely transferred • More electronegative atom: δ- . (δ represents the partial negative charge formed) • Less electronegative atom: δ+

16. Lewis Structures • represents how an atom’s valence electrons are distributed in a molecule • Show the bonding involves (the maximum bonds can be made) • Try to achieve the noble gas configuration

17. Rules • Duet Rule: sharing of 2 electrons • E.g H2 • H : H • Octet Rule: sharing of 8 electrons • Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule • E.g F2, O2 • Bonding pair: two of which are shared with other atoms • Lone pair or nonbonding pair: those that are not used for bonding

18. 1 Lewis Structures of Molecules with Multiple Bonds • Use 6N + 2 Rule • N = number of atoms other than Hydrogen • If • Total valence – (6N + 2) = 2 •  1 double bond • Total valance e- - (6N + 2) = 4 •  two double bonds or 1 triple bond

19. Examples • Write the Lewis structure of CH3F, ClO3-, F2

20. 1.7 Formal Charges • Difference between the number of outer-shell electrons “owned” by a neutral free atom and the same atom in a compound

21. Examples • Determine the formal charge for each atom in the following molecules • NH4+ • NO2- • CO32-

22. Resonance • Whenever a molecule or ion can be represented by two or more Lewis structures that differ only in the position of the electrons • None of these resonance structures will be a correct representation for the molecule or ion • The actual molecule or ion will be better represented by a hybrid or hypothetical structures • Represented by a double headed arrows ( )

23. Examples

24. Resonance - stabilization • The more covalent bonds a structure has, the more stable it is

25. Resonance-stabilization • Structure in which all the atoms have a complete valence shell of electrons are especially

26. Resonance stabilization • Charge separation decrease stabilization • Resonance contributors with negative charge on highly electronegative atoms are stable ones with negative charge on less or nonelectronegative atoms

27. 1.9 Quantum Mechanisms and Atomic Structure • Schröndinger’s quantum mechanical model of atomic structure is frame in the form of a wave equation; describe the motion of ordinary waves in fluids. • i. Wave functions or orbitals (Greek, psi , the mathematical tool that quantum mechanic uses to describe any physical system • ii. 2 gives the probability of finding an electron within a given region in space • iii. Contains information about an electron’s position in 3-D space • defines a volume of space around the nucleus where there is a high probability of finding an electron • say nothing about the electron’s path or movement

28. 11.2 Electromagnetic Radiation • Radiation energy – has wavelike properties • Frequency (υ, Greek nu) – the number of peaks (maxima) that pass by a fixed point per unit time (s-1 or Hz) • Wavelength (λ, Greek lambda) – the length from one wave maximum to the next • Amplitude – the height measured from the middle point between peak and trough (maximum and minimum) • Intensity of radiant energy is proportional to amplitude

29. Wave function

30. Heisenberg Uncertainty Principle – both the position (Δx) and the momentum (Δmv) of an electron cannot be known beyond a certain level of precision 1. (Δx) (Δmv) >h 4π 2. Cannot know both the position and the momentum of an electron with a high degree of certainty 1.10 Atomic orbital

31. Molecular Orbitals • Two types of atomic of atomic orbitals are combined as they come close to each other • Hybridization: blending combination of atomic orbitals to form new orbital • Carbon has three possible molecular orbitals • sp3 sp2 sp

32. Orbitalsrepsonsible for creating the covalent bonds • 2 special names for covalent bonds of organic molecules

33. sp3 orbitals responsible for creating all “single bonds” of all organic molecules  alkanes sp3 molecular orbitals

34. Examples

35. sp2 molecular orbitals • All sp2 molecular orbitals responsible for creating all double bonds in organic molecules •  alkenes

36. Examples

37. 1.13B – Cis –Trans Isomerism • Which of the following alkene can exist as cis-trans isomers? Write their structure

38. sp molecular orbitals • All sp orbitals responsible for creating all triple bonds of organic molecules •  alkynes

39. Examples

40. Examples • Draw a bonding picture for the following molecule, showing all π, σ – bonds using σ-framework and π-framework

41. Molecular Orbitals • Two types: • Bonding molecular orbitals • Contains both electrons in the lowest energry state or ground state • Formed by intereaction of orbitals with same phase signs • Increases the propability

42. Molecular orbitals • Antimolecular orbitals • Contains no electrons in the ground state • Formed by intereaction of orbitals with opposite phase signs • Result with nodes

43. Molecular orbitals

44. Shape of Molecules • VSEPR Theory • Valence shell electron pair repulsion • Bond angles and geometry • Steric number = # bond to - # lone pairs central atom to central atom • Rules: 1- Carbon will always be the central atom 2 – Double bond; triple bonds will count as 1 bond

45. Shape of Molecules

46. Molecular shapes • VSEPR method can be used to predict the shapes of molecules containing multiple bonds • Assume that all electrons of a multiple bond act as one unit

47. Examples • Use VSEPR theory to predict the geometry of each of the following molecules and ions • SiF4 • BeF2

48. 1.17 Representation of Structural formulas • Structual formula for propyl alcohol

49. Dash Structure • Atoms are joined by single bonds can rotate relatively freely with respect to one another

50. Dash Formula - Isomerism