1. The Basic – Bonding and Molecular Structure
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. 1.10 Atomic orbital
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

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. Orbitals repsonsible for creating the covalent bonds
2 special names for covalent bonds of organic molecules
Sigma (σ) bond
Pi (π) bond
Created when “head on” overlap occurs of orbitals
Created when “side on” overlap occurs of orbitals

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

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

51. Condensed Structural Formulas
All hydrogen atoms are written immediately after the carbon that they’re attached

52. Bond-Line Formulas
Hydrogen and carbon atoms will not appear in the formula
Each end of the line represents carbon atom

53. Examples
For each of the following, write a bond line formula