Organic chemistry A Chapter 1 Introduction By Prof. Dr. Adel M. Awadallah

1. Organic chemistry A Chapter 1 Introduction By Prof. Dr. Adel M. Awadallah Islamic University of Gaza

2. Chapter 1 Bonding and Isomerism Organic chemistry is the chemistry of the compounds of carbon Atoms consist mainly from • Nucleus: (containing Protons and Neutrons) Protons (positive particles, Atomic Number) Neutrons (Neutral particles) Protons + Neutrons (Atomic weight) b) Electrons: negatively charged particles

3. Electronic Configuration (review) • Shells (n = 1,2,3,4, ….) • (Subshells , s, p, d, f) • (orbitals: rgions of space around the nucleus containing electrons) • each orbital contains only 2 electrons with different spins

5. Ionic Compounds They are formed by the transfer of one or more valence electrons from one atom to another Electropositive atoms: give up electrons and form cations. Electronegative atoms: accept electrons and form anions Ionic compounds: are composed of positively charged cations and negatively charged cations

6. Covalent compounds A covalent bond: is formed when two atoms share one or more electron pairs. A molecule consists of two or more atoms joined by covalent bonds Bond energy: is the energy necessary to break a mole (6.022 x 1023)of covalent bonds. Bond length: is the average distance between two covalently bonded atoms

7. Valence and bonding in organic compounds

8. Polarity of Bonds: depends on the electronegativity difference

9. Examples:

10. Polarity of Molecules: depends on the sum of the polarity of bonds (geometry)

11. Isomers: different compounds having the same molecular formula

12. Abbreviated structural formulas

13. Formal Charges: are the charges that each atom carries, and can be calculated as follows • Formal charge = Valence electrons – bonds – electrons • Example: • Resonance: arises whenever we can write two or more structures for a molecule with different arrangements of electrons but identical arrangement of atoms

14. Arrows:

15. The orbital view of bonding

16. Orbital picture of Methane and ethane The bond formed by end-to-end overlap is called a sigma bond.

17. Bond Angles in Methane

18. Bonding in Ethylene (ethene)

19. A pi bond is one in which the electrons in the p orbitals are held above and below the plane of the molecule. The sigma bond is stronger than the pi bond. A double bond is formed from a sigma bond and a pi bond, and so it is stronger than a single bond.

20. Bonding in acetylene (ethyne)

21. Classification of organic compounds according to functional groups

22. Intramolecular Forces Bond Dissociation Energy:The amount of energy consumed or liberated when a bond is broken orformed

25. Intermolecular Forces Intermolecular forces:attractive forces between molecules. • Intramolecular forces: hold atoms together in a molecule.(chemical bonding) • Intramolecular forces stabilize individual molecules, whereas intermolecular forces are primarily responsible for the bulk properties of matter (i.e. melting point and boiling point). • Intermolecular vs Intramolecular • 41 kJ to vaporize 1 mole of water (inter) • 930 kJ to break all O-H bonds in 1 mole of water (intra) • Generally, intermolecular forces are much weaker than intramolecular forces. • The boiling points of substances often reflect the strength of the intermolecular forces operating among the molecules. At the boiling point, enough energy must be supplied to overcome the attractive intermolecular forces between liquid molecules. The same is true for melting a solid

26. Types of Intermolecular forces: • 1- Dipole-Dipole Forces:are attractive forces that act between polar molecules • The larger the dipole moments, the greater the force. • The Hydrogen Bond:is a special type of dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. • Aــــ H---B or Aــــ H---A • A & B are N, O, or F • The average energy of a hydrogen bond is quite large for a dipole-dipole interaction (up tp 40KJ/mol). Thus, hydrogen bonds are a powerful force in determining the structure and properties of many compounds.

27. 2- Ion-Dipole Forces:Attractive forces between an ion (either a cation or an anion) and a polar molecule. • 3- Dispersion (London) force:arise as a result of temporary dipoles induced in the molecules or atoms. • - act between all molecules. • - only force between nonpolar molecules and noble gas atoms. • - strength depends primarily on size of molecule, very weak for small molecule but fairly strong for large ones. (Shape plays a role too.)

29. Melting points • Melting is the change from the highly ordered arrangement of particles in the crystalline lattice to the more random arrangement that characterizes a liquid • Mp NaCl = 801 • Mp CH4 = -183

30. Boiling point:The temperature at which the vapor pressure of a liquid equals the external pressure Boiling involves the breaking away from the liquid of individual molecules or pairs of oppositely charged ions Associated liquids (Hydrogen bonded liquids) Have relatively high bp. HF boils 100 oC higher than HCl H2O boils 160 oC higher than H2S

31. Solubility:Like dissolve like Ionic Solutesare dissolved in water or very polar solvents. They form ion-dipole bonds Non-ionic solutes: solubility depends on polarity Methane (London forces ) in CCl4 (London forces ) Methanol (hydrogen bonding) in water (hydrogen bonding)

32. Acids and bases: Arrhenius DefinitionAcid: H+ donor, Base: OH- donor HCl + NaOH → NaCl + H2O Lowry-Brønsted Acids and Bases: Acid: H+ donor , Base: H+ acceptor HCl + :OH- → Cl- + H2O HCl + :NH3 → Cl- + NH4+ Lewis Definition: Acid: e acceptor, base:electron donor H+ + :OH-→ H2O F3B + :NH3 → F3B:NH3