Organic chemistry A Chapter 1 Introduction By Prof

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

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

Electronic Configuration (review) 1) Shells (n = 1, 2, 3, 4, …. ) 2) (Subshells , s, p, d, f) 3) (orbitals: rgions of space around the nucleus containing electrons) 4) each orbital contains only 2 electrons with different spins Subshell Number of orbitals (electrons) Name of orbitals S P 1 (2) 3 (6) s Px, py, pz d 5 (10) Dxy, dyz, dxz, dz 2, 2 2 d. X -y

Examples: Filled Shells Play almost no role in chemical reaction Valence Electrons They are the outermost electrons and they are mainly involved in chemical bonding

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

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

• Valence and bonding in organic compounds

• Polarity of Bonds: depends on the electronegativity difference

• Examples:

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

• Isomers: different compounds having the same molecular formula

• Abbreviated structural formulas

• 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

• Arrows:

• The orbital view of bonding

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

Bond Angles in Methane

• Bonding in Ethylene (ethene)

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 a pi bond, and so it is stronger than a single bond.

• Bonding in acetylene (ethyne)

• Classification of organic compounds according to functional groups

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

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 k. J to vaporize 1 mole of water (inter) • 930 k. J 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

• 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 an electronegative O, N, or F atom. A ــــ H---B or A ــــ 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 40 KJ/mol). Thus, hydrogen bonds are a powerful force in determining the structure and properties of many compounds.

• 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. )

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 Na. Cl = 801 • Mp CH 4 = -183 •

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 o. C higher than HCl H 2 O boils 160 o. C higher than H 2 S

Solubility: Like dissolve like Ionic Solutes are dissolved in water or very polar solvents. They form ion-dipole bonds Non-ionic solutes: solubility depends on polarity Methane (London forces ) in CCl 4 (London forces ) Methanol (hydrogen bonding) in water (hydrogen bonding)

Acids and bases: Arrhenius Definition Acid: H+ donor, Base: OH- donor HCl + Na. OH → Na. Cl + H 2 O Lowry-Brønsted Acids and Bases: Acid: H+ donor , Base: H+ acceptor HCl + + : OH: NH 3 → → Cl. Cl- + + H 2 O NH 4+ Lewis Definition: Acid: e acceptor, base: electron donor H+ + : OH→ H 2 O F 3 B + : NH 3 → F 3 B: NH 3