1. Five Traditional Branches of Chemistry:
Organic
Analytical
Inorganic
Physical
Biochemistry

2. Organic Chemistry references
Organic Chemistry references
Morrison & Boyd (6th)
Organic Chemistry by:
William H. Brown
Christopher S. Foote
Brent L. Iverson

3. Organic Chemistry
The study of the compounds of carbon, their properties and the changes that they undergo. 
Is the chemistry of compounds that contain the element carbon.
Include description of
(1- Nomenclature 2- Synthesis 3- Reactions 4- Mechanisms)

4. C is a small atom
it forms single, double, and triple bonds
it is intermediate in electronegativity (2.5)
it forms strong bonds with C, H, O, N, and some metals

5. Crude
oil
Coal

6. Composition of matter
All matter is composed of the same building blocks called atoms. There are two main components of an atom.
Nucleus contains positively charged protons and uncharged neutrons. Most of the mass of the atom is contained in the nucleus.
Electron cloud is composed of negatively charged electrons. The electron cloud comprises most of the volume of the atom.

7. neutral atom
Atoms are electrically neutral because they have equal number of proton (p) and electron (e).
The charge on a proton is equal in magnitude but opposite in sign to the charge on an electron.
Ex; Hydrogen has (1p and 1e ) and they cancel each other.
The number of protons in the nucleus equals the number of electrons. This quantity, called the atomic number, is unique to a particular element.
For example, every neutral carbon atom has an atomic number of six, meaning it has six protons in its nucleus and six electrons surrounding the nucleus.

9. Ion Atoms can gain or lose electrons when they form ionic compounds.
When atoms lose or gain electrons they become charged Atoms having a (+) or (–) charges are called ions .
A cation is positively charged and has fewer electrons than its neutral form.
An anion is negatively charged and has more electrons than its neutral form.
Sodium (Na) loses an electrons when it forms NaCl and become Na+.

10. Formation of Ions:
We use a single-headed curved arrow to show the transfer of one electron from (Na) to (F).
In forming ( Na+ F-) the single (3s) electron from (Na) is transferred to the partially filled valence shell of (F).

11. Isotopes
Are two atoms of the same element having a different number of neutrons.
Most carbon atoms have six protons and six neutrons in the nucleus, but 1.1% have six protons and seven neutrons.

12. Bonding Bonding is the joining of two atoms in a stable arrangement.
Hydrogen gas (H2), formed by joining two hydrogen atoms,
Methane (CH4), the simplest organic compound, formed by joining a carbon atom with four hydrogen atoms.
Atoms attain a complete outer shell of valence electrons.
Because the noble gases in column 8 of the periodic table are especially stable as atoms having a filled shell of valence electrons, the general rule can be restated.

13. A first-row element like hydrogen can accommodate two electrons around it.
This would make it like the noble gas helium at the end of the same row.
A second-row element is most stable with eight valence electrons around it like neon.
Elements that behave in this manner are said to follow the octet rule.

14.  Ionic bond:
Ionic bond formed by the transfer of valence electrons from one element to another to achieve noble gas electron configurations, result ions held together by electrostatic attraction.

15. The simplest covalent bond is H2
Covalent bond formed by the sharing of valence electrons to achieve noble gas electron configurations.
The simplest covalent bond is H2
the single electrons from each atom combine to form an electron pair.
The number of shared pairs:
one shared pair forms a single bond
two shared pairs form a double bond
three shared pairs form a triple bond

16. Polar and Non-polar Covalent Bonds:
Although all covalent bonds involve sharing of electrons, they differ widely in the degree of sharing We divide covalent bonds into
Non-polar covalent bonds
Polar covalent bonds

17. Example of a polar covalent bond is that of (H-Cl)
we show polarity by using the symbols d+ and d-, or by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end.

18. Bond dipole moment (m):
Polar Covalent Bonds
Bond dipole moment (m):

19. Fluorine F (at.no 9) 9p=9e 1s2,2s2,2p5 Chlorine Cl (at.no 17) 17p=17e
  1s2,2s2,2p6,3s2,3p5
 Bromine Br (at.no 35) 35p=35e
1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p5
 Iodine I (at.no 53) 53p=53e
  1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p6,5s2,4d10,5p5
order of filling

20. Hybridization of atomic orbitals:
sp = linear; 180o
sp2 = trigonal; 120o
sp3 = tetrahedral; 109.5o

21. Hybridizations of carbon

23. Polarity:
Covalent bonds are polar when the two atoms sharing electrons have different electronegativities.
eg. H—Cl
δ+ δ-
A charge separation or a dipole gives a polar bond.
O :O=O: has a non-polar bond

24. Electronegativity
Is the measure of attraction of atom’s for the electrons it shares with another atom in a chemical bond.

25. Intramolecular forces. Attractions between atoms.
Ionic Attractions (very strong) Na+ Cl-
dipole-dipole attractions H—Br
hydrogen bonding ( H attached to N,O,F )
(H—O------H—O)
van der Waals (London forces) (weak) Br—Br

27. Solubility “like dissolves like”
~ water soluble? must be ionic or highly polar + H-bond
(hydrophilic)
~ water insoluble? must be non-polar or weakly polar
(hydrophobic)

29. Within the period of the periodic table, acid strength increases with increasing electronegativity:
CH4 < NH3 < H2O < HF
Within a group of elements, acid strength increases with increasing size:
HF < HCl < HBr < HI

30. F- > Cl- > Br- > I-
Which one is the stronger acid?
H2O or H2S?
oxygen & sulfur are in the same group and sulfur is bigger in size: H2S > H2O
What is the order of base strength?
F Cl Br I-
in the halogen family base strength decreases with increasing size:
F- > Cl- > Br- > I-