Five Traditional Branches of Chemistry:

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

2. Organic Chemistry • Organic Chemistry references • Organic Chemistry -3rd Edition-Janice Gorzynski Smith • 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- Nomenclature2- 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. Nucleuscontains positively charged protonsand 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 to the charge on an electron but opposite in sign and magnitude. • 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. • 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 thetransfer of valence electrons from one element to another to achieve noble gas electron configurations, result ions held together by electrostatic attraction.

15. Covalent bond: • 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. Polar Covalent Bonds Bond dipole moment (m):

19. van der Waals Forces • Also called London forces, are very weak interactions caused by the momentary changes in electron density in a molecule. van der Waals forces are the only attractive forces present in nonpolar compounds. • For example, although a nonpolar CH4 molecule has no net dipole, its electron density may not be completely symmetrical, creating a temporary dipole.

20. This can induce a temporary dipole in another CH4 molecule, with the partial positive and negative charges arranged close to each other. • The weak interaction of these temporary dipoles constitutes van der Waals forces. • The larger the surface area, the larger the attractive force between two molecules, and the stronger the intermolecular forces. • molecules such as CH3CH2CH3(propane) have stronger van der Waals interactions than CH4(methane).

22. 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

23. Sp3 Hybridization Of Carbon

24. Hybridizations of carbon

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

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

28. 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

31. Within the periodof the periodic table, acid strength increases with increasingelectronegativity: • CH4 < NH3 < H2O < HF • Within a groupof elements, acid strength increases with increasingsize: • HF < HCl < HBr < HI

32. 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-