1. Chapter 12. Saturated Hydrocarbons
Sections

2. 12. 1 Organic and Inorganic Compounds 12
12.1 Organic and Inorganic Compounds 12.2 Bonding Characteristics of the Carbon Atom 12.3 Hydrocarbons and Hydrocarbon Derivatives 12.4 Alkanes: Acyclic Saturated Hydrocarbons 12.5 Structural Formulas 12.6 Alkane Isomerism 12.7 Conformations of Alkanes 12.8 IUPAC Nomenclature for Alkanes 12.9 Line-Angle Formulas for Alkanes Classification of Carbon Atoms Branched-Chain Alkyl Groups Cycloalkanes IUPAC Nomenclature for Cycloalkanes Isomerism in Cycloalkanes Sources of Alkanes and Cycloalkanes Physical Properties of Alkanes and Cycloalkanes Chemical Properties of Alkanes and Cycloalkanes Chemistry at a Glance: Properties of Alkanes and Cycloalkanes

3. Introduction
Organic chemistry is the study of the compounds of carbon atoms consisting of single, double, and triple covalent bonds with C, H, O, N, and some metals
Biochemistry is the study of the compounds of carbon atoms consisting of single, double, and triple covalent bonds with C, H, O, N, and some metals

4. 12.1 Organic and Inorganic Compounds

5. Sections Chapter 1: Covalent Bonding and Shapes of Molecules
Introduction
Electronic Structure of atoms
Lewis Model of bonding
Bond angles and molecular shapes
Polar and non-polar molecules
Resonance
Orbital overlap of covalent bonding
Functional groups

6. Introduction
Organic chemistry is the study of the compounds of carbon.
C is a small atom
it forms four bonds consisting of single, double, and triple bonds
it forms strong bonds with C, H, O, N, and some metals
Problem: Identify organic and inorganic compounds:
NaCl, C2H5OH, CH3COOH, Na2CO3, CH3OK, KOH

7. Atomic Orbitals
s orbital - a spherical-shaped atomic orbital; can hold a maximum of 2 electrons
p orbital - a dumbbell-shaped atomic orbital; the three p orbitals (px, py, pz) can hold a maximum of 2 electrons each
Electrons always fill starting with the lowest-energy orbital:
lower energy higher energy
1s2 2s2 2p6 3s2 3p6
We will be concerned with only the valence electrons which are the outermost electrons involved in forming bonds.

8. Shapes of s Atomic Orbitals
All s orbitals have the shape of a sphere, with its center at the nucleus
of the s orbitals, a 1s orbital is the smallest, a 2s orbital is larger, and a 3s orbital is larger still

9. Shapes of p Atomic Orbitals
A p orbital consists of two lobes arranged in a straight line with the center at the nucleus

10. Electronic Structure of atoms (Review)
Ground state electronic configuration of atoms in core format
Carbon (C): ): [He] 2s2, 2p2
or [He] 2s2, 2px13py13pz0
Potassium (K): {Ar] 4s1
Phosphorous (P): [Ne] 3s2, 3p3
Valence shell electronic configuration
Carbon (C): ): 3s2, 3p2
Potassium (K): 4s1
Phosphorous (P): 3s2, 3p3
How you get the electronic configuration of an atom from the periodic table?

11. Excited State Valence Electron Configuration
Carbon (C):
Ground state: 2s2, 2p2 or 2s2, 2px13py13pz0
Excited State: 2s1, 2p3 or 2s1, 2px13py13pz1

12. Lewis Model of bonding (Review)
"octet rule“
atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell
“Lewis structure of atoms”
Shows only valence electrons, is a convenient way of representing atoms to show their chemical bonding pattern.

13. Lewis structure of atoms (Review)

14. Cations and Anions Cations Anions C N Na+ O
What elements lose electrons? And how many?
What is the positive charge on their cations?
Anions
What elements gain electrons?
What is the positive charge on their anions?
Covalent bonds
How many covalent bonds are formed?
What elements share electrons? C N
Na+ O

15. Covalent or Ionic Identify covalent and ionic compounds:
NaCl, C2H5OH, CH3COOH, Na2CO3, CH3OK, KOH
Covalent :
Ionic:

16. Ionic model of bonding model (Review)
Ionic bond - results from the electrostatic attraction between a cation and an anion of two atoms typically involves a metal and a nonmetallic element.
Anion: An atom that gains electrons becomes a negative ion
Cation: An atom that loses electrons becomes a positive ion

17. Covalent model of bonding (Review)
Covalent bonds - results from the sharing of electrons between two atoms typically involves two nonmetallic elements

18. Electronegativity (Review)
Is the attraction of an atom for its valence electrons H
2.1
Li Be B C N O F
Na Mg Al Si P S Cl
increases i n c r e a s

19. Nonpolar/polar-covalent and ionic bonds. (Review)
We classify chemical bonds as polar covalent, nonpolar covalent and ionic based on the difference in electronegativity between the atoms

20. Classify following bonds nonpolar-covalent, polar-covalent or ionic bonds
N-H nonpolar-covalent, polar-covalent or ionic bonds
O-H nonpolar-covalent, polar-covalent or ionic bonds
C-H nonpolar-covalent, polar-covalent or ionic bonds
C-F nonpolar-covalent, polar-covalent or ionic bonds
Na-Cl nonpolar-covalent, polar-covalent or ionic bonds
Al-Cl nonpolar-covalent, polar-covalent or ionic bonds

21. Drawing Lewis structure molecules and ions (Review)
1) Predict arrangement of atoms.
H is always a terminal atom.
Halogens and oxygen are often terminal.
The central atom of binary compounds is usually written
first and has the lowest subscript.
Most organic compounds have more than two central atoms.
These are mainly C, but N, O and S can also be central atoms.
2) Total number of valence electrons (e-)
Add all valence electron of atoms in the molecule from the formula.
Add the ion charge for negative ions or subtract for positive ions.

22. Drawing Lewis structure molecules and ions (Review)
3) Draw the skeletal structure by connecting the atoms with single bonds.
4) Give each of the atoms an octet (8 e-). Adding unshared pairs of electrons
5) Count the total number of e- used through step 4 and compare to the number calculated in step 2.
a) If it results in zero, the structure is correct.
b) For every two electrons too many, another bond is added (minimize formal charges).
Multiple bonds form only with C, N, O and S.
Total number of bonds to neutral atoms:
4 bonds to C
3 bonds to N, P
2 bonds to O, S
1 bond to H, F, Cl, Br, I

23. Calculation of formal charges of atoms in the Lewis structure
1. For a neutral molecule, the sum of the formal charges equals zero. For a polyatomic ion, the sum of the formal charges equals the charge on the ion.
2. Formal charge of each atom is calculated by:
(group #) - (# unshared e-) - ½ (# shared e-)
3. Formal charges are shown as + or - on the atom with that charge.
4. An atom with the same number of bonds as its group number has no formal charge.
5. In a molecule if two different elements can be assigned a negative charge, then the more electronegative element gets the charge; the same sign should not be given to bonded atoms.

24. H Cl Types of electrons Bonding pairs
Two electrons that are shared between two atoms.
A covalent bond.
Unshared (nonbonding ) pairs
A pair of electrons that are not shared between two atoms. Lone pairs or nonbonding electrons. oo
Unshared
pair
H Cl oo oo oo
Bonding pair

25. Lewis Structures (Review)
How many bonding electron pairs are in the molecule?
How many bonding electron pairs are in each atom?
How many nonbonding electron pairs are in the molecule?

26. Draw Lewis structure of molecules
CHCl3
C2H4
C3H8O
CH3CH2CH2OH
CH3CH2OCH3
CH3CO2H
CH3CHO

27. Draw Lewis structure and assign formal charges
CH3NH3+
CH3O-

28. Valence-Shell Electron-Pair Repulsion (VSEPR) model (Review)
For predicting shapes of molecules and polyatomic Ions based on the repulsion of valence pairs of electrons making them as far apart as possible around an atom of a Lewis structure.
1) Draw the Lewis structure for the molecule or ion.
2) Determine the number of bonding and unshared pairs attached to the central atom.
One single, double or triple bond counted as a bonding pair
3) Choose the appropriate case from the given chart.

29. Predict the bond angles of molecules from their Lewis structures
Predict the bond angles of molecules from their Lewis structures. (Review)
Molecule
Bonding pairs
unshared pairs
Shape
H2O
two
bent
NH3
three
one
Trigonal pyramid
CH2O
none
Trigonal planar
CHCl3
four
tetrahedral
CH4
CCl4

30. Polar and nonpolar molecules

31. Molecular Shape and Polarity (Review)
Molecule
Bonding pairs and unshared pairs
Electron pair distribution
Polarity
H2O
four
asymmetric
polar
NH3
CH2O
three
CHCl3
CH4
symmetric
Non-polar
CCl4

32. Curved arrow Electron pushing
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 an atom to an adjacent bond
Electron pushing by the use of curved arrows is also used in explaining reaction mechanisms

33. Drawing Curved Arrows Arrows can be drawn from: 1) lone pair 2) bond
To show the movement of electrons in breaking and forming bonds. The tail of the arrow is started at the site of electron density (negative character such as a pi bond or lone pair of electrons) and proceeds to the arrowhead which is drawn to the site of electron deficiency (positive character).
NEGATIVE TO POSITIVE!
Arrows can be drawn from:
1) lone pair
bond
lone pair
2) bond
bond
3) bond

34. Resonance
For many molecules and ions with double bonds, two or more Lewis structure could be written

35. Orbital Overlap Model

36. Hybrid Atomic Orbitals
Hybridization is the mixing up of two or more atomic orbitals
There are three types of hybrid atomic orbitals for carbon
sp3 (one s orbital + three p orbitals give four sp3 orbitals)
sp2 (one s orbital + two p orbitals give three sp2 orbitals)
sp (one s orbital + one p orbital give two sp orbitals)

37. s and p hybrids
Four sp3 hybrids
Three sp2 hybrids
Two sp hybrids

38.  and  bonds
Overlap of hybrid orbitals can form two types of bonds, depending on the geometry of the overlap
bonds are formed by “direct” overlap
 bonds are formed by “parallel” overlap of unhybrid p prbitlas

39.  and  bonds in single and multiple bonds
single bond - one shared pair of electrons between two atoms; a s bond
double bond - two shared pairs of electrons between two atoms; one s bond and one p bond
triple bond - three shared pairs of electrons between two atoms; one s bond and two p bonds

40. Bond Properties Bond strength: strongest weakest Bond length:
longest shortest

41. Counting  and  bonds in Lewis structure

42. Predicting hybridization of atoms in a Lewis structure
Count sigma bonds and unshared electrons around the atom
If the total number of pairs:
2 sp hybridization
3 sp2 hybridization
4 sp3 hybridization

43. Functional Groups in Organic Compounds
Functional group: an atom or group of atoms within a molecule that shows a characteristic set of physical and chemical properties
Functional group
divide organic compounds into classes
the sites of characteristic chemical reactions
the basis for naming organic compounds

44. Common Functional Groups
Alcohol(carbon, hydrogen and oxygen)
CH3-CH2-OH
Thiol (thioalcohol)   thio means sulfur
CH3-CH2-SH
Amine
CH3-CH2-NH2
Ether
CH3-CH2-O-CH2-CH3

45. Common Functional Groups (continued)
Carbonyl group
or unit
           
Aldehyde
         
ketones
        
Carboxylic acids
   

46. Classification of organic compounds
Class Functional group Example